chap17

2014-10-01
ELECTROCHEMISTRY
17
CHAPTER
17.1 Electrochemical Cells
17.2 Cell Potentials and the Gibbs Free Energy
17.3 Molecular Interpretation of Electrochemical
Processes
17.4 Concentration Effects and the Nernst Equation
17.5 Molecular Electrochemistry
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17.1 ELECTROCHEMICAL CELLS
 Electrochemical reactions interconvert chemical and electrical
energy through the oxidation (anode) and reduction (cathode)
half reactions occurring on the surfaces of electrodes.
 Galvanic cell (or Voltaic cell):
~ Spontaneous chemical reaction (G < 0)  producing electricity
 Electrolytic cell
~ External input of electricity
 driving nonspontaneous chemical reactions (G > 0)
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 Galvanic Cells
Fig. 17.1 & 2 Cu (anode) – Ag (cathode) in a Galvanic cell.
Redox reaction in the above cell:
Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)
cathode (+) reduction half reaction: Ag+(aq) + e–  Ag(s)
anode (–) oxidation half reaction: Cu(s)  Cu2+(aq) + 2 e–
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 Line Notation:
Cu(s) | Cu(NO3)2(aq) || AgNO3(aq) | Ag(s)
or
Cu | Cu2+ || Ag+ | Ag
| phase boundary,
|| salt bridge
left: anode (negative), right: cathode (positive)
Commas are used to indicate the presence of different
components in the same phase: | H+,Cl– |
Half cells are physically separated but connected by a salt bridge.
~ Allows ion migration in solution but prevents extensive mixing
of electrolytes
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 Electrolytic Cells
Galvanic cell
Fig. 17.3 A copper-silver electrolytic cell.
 Electrolytic cell:
Reversing the direction of a spontaneous process requires the
application of an external potential (a power supply or a battery).
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 The Coulomb (C) is defined as the total charge transferred
by 1 A of current flowing for 1 s: 1 C = (1 A)(1 s)
 Total charge transferred by i A of current in t seconds
= Total charge carried by n moles of electrons
Q = it = nF
n: number of moles of electrons
 Equivalent mass in a redox reaction: Weq = M / n’
M: molar mass
n’: number of moles of electrons transferred per mole
of substance in the corresponding redox reaction
Ex. Zinc-silver Galvanic cell
(cathode) Ag+(aq) + 1e–  Ag(s) Weq(Ag) = (107.8)/1 = 107.8 g mol–1
(anode) Zn(s)  Zn2+(aq) + 2e– Weq(Zn) = (65.38)/2 = 32.69 g mol–1
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 Faraday's Law (1833)
1. The mass (W) of a substance that is produced or consumed
in an electrochemical reaction is proportional to the quantity
of electric charge passed. W  Q  n
2. Equivalent masses (Weq) of different substances are produced
or consumed in electrochemical reactions by a given amount
of electric charge passed. W  Weq  (n’)–1
1+ 2 W = nWeq
 Faraday Constant, F: F = e NA = 96,485 C mol–1
The charge of 1 mole of electrons.
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Cu-(AgCl/Ag) Electrolytic cell. 0.500 A current for 101 min.
Mass of Cu dissolved and the mass of Ag deposited?
EXAMPLE 17.2
1
3
it  0.500 C s  6.06 10 s 
n 
 3.14 102 mol e
1
F
96, 485 C mol
Half-reactions:
(cathode) AgCl(s) + 1e–  Ag(s) + Cl–(aq)
(n’ = 1)
Weq(Ag) = (107.8)/1 =107.8 g mol–1
(anode) Cu(s)  Cu2+(aq) + 2e–
(n’ = 2)
Weq(Cu) = (63.55)/2 = 31.78 g mol–1
W(Ag) = nWeq(Ag) = (3.14ⅹ10–2 mol)(107.8 g mol–1) = 3.38 g
W(Cu) = nWeq(Cu) = (3.14ⅹ10–2 mol)(31.78 g mol–1) = 0.998 g
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17.2 CELL POTENTIALS AND THE GIBBS FREE
ENERGY
 Electrical work, welec
~ Change in the potential energy, EP (in joules), associated
with the transfer of Q coulombs of negative charge through
a potential difference E (in volts) given by
welec = EP = – Q E  – QEcell
= – Q (Ecathode – Eanode) = – itEcell
Ecell > 0 for a galvanic cell
~ work done by the cell producing electrical work
Ecell < 0 for an electrolytic cell
~ work done on the cell by an external power supply
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 Maximum electrical work utilizing a spontaneous reaction
welec,rev  welec,max  G (at constant T and P )
(Derivation)
G = H – TS = E + PV – TS
G = E + PV – TS (const T and P)
= q + welec – Pext V + PV – TS
(E = q + w, w = welec – Pext V )
For a reversible process, Pext = P and qrev = TS.
welec,rev = G
If n moles of electrons pass through the external circuit
of a reversible galvanic cell,
welec,rev  G  QEcell  nFEcell
(reversible)
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 Standard Reduction Potentials
 Standard hydrogen electrode (SHE)
~ Primary reference electrode
2 H3O+(aq) + 2 e–
 H2(g) + 2 H2O(l), Eo(SHE) = 0.00 V
 H3O    1 M (or a H O  1)
3
pH2  1 atm (or a H 2  1)
 Standard cell potentials of half reactions
are measured with reference to the SHE.
 Standard cell potential of a Galvanic cell
Fig. 17.4 Schematic of a
standard hydrogen electrode
o
o
o
Ecell
 Ecathode
 Eanode
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Ex. 2 Calculate the standard cell potential for the Cu|Cu2+||Ag+|Ag cell.
Ag+(aq) + e–  Ag(s),
Eo = 0.799 V
Cu2+(aq) + 2 e–  Cu(s),
Eo = 0.340 V
Cathode:
2 Ag+(aq) + 2 e–  2 Ag(s)
Anode:
Cu(s)  Cu2+(aq) + 2 e–
Oveall:
2 Ag(s) + Cu2+(aq)  2 Ag+(aq) + Cu(s)
o
o
o
 Ecathode
 Eanode
 0.799 V  (0.34 V)  0.459 V
 Cell potential: Ecell
*The standard cell potential is an intensive property.
 Gibbs free energy change (or work done by the system)
o
G o   nFEcell
 ( 2 mol)(96,500 C mol 1 )(0.459 V)  0.886 kJ
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 Adding and subtracting half-cell reactions
Ex. 3 Find Eo for a half-reaction from Eo’s of other two half-reactions.
Cu2+(aq) + 2 e–  Cu(s),
E1o = Eo (Cu2+|Cu) = 0.340 V
Cu+(aq) + e–  Cu(s),
E2o = Eo (Cu+|Cu) = 0.522 V
Cu2+(aq) + e–  Cu+(aq),
E3o = Eo (Cu2+|Cu+) = ?
G (not Eo ) is an extensive property.
Ghco = –nhc FEo
G3o = G1o – G2o = –n1FE1o + n2FE2o = –n3FE3o
E3o = (n1 E1o – n2 E2o) / n3 = (2  0.340 – 1  0.522) / 1 = 0.158 V
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 Oxidizing and Reducing Agents
 Oxidizing agent: easily reduced, large positive Eo (F2, H2O2, MnO4–)
 Reducing agent: easily oxidized, large negative Eo,
(alkali & alkaline earth metals)
 Oxidizing powers of O2 and O3
~ Effective oxidizing agents in acidic solution at pH 0:
O2 + 4 H 3 O+ + 4 e –  6 H 2 O
Eo = 1.229 V
O3 + 2 H3O+ + 2 e–  O2 +3 H2O
Eo = 2.07 V
O3 is a stronger oxidizing agent than O2 because Gfo (O3 )  0 .
 G for the reduction of H3O+(aq) by O3 is more negative than
G for its reduction by O2.
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 Reduction Potential Diagram
 Latimer diagram:
A species can disproportionate if and only if
Eo (left; reverse) < Eo (right; forward).
Then, thermodynamically feasible disproportionation of Cu+ is
2 Cu+  Cu2+ + Cu
Eo = 0.522 V (right) – 0.158 V (left) = + 0.364 V > 0
Eo > 0  Go < 0 , spontaneous process !
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 Alternative Reference Electrodes
 Saturated calomel electrode
Hg2Cl2 (calomel) + 2 e–
 2 Hg + 2 Cl–(saturated)
Eo = 0.242 V
~ Phased out due to environmental
problem of Hg contamination
 Ag/AgCl electrode or
Ag|AgCl|KCl(saturated, aq) electrode
AgCl + e–  Ag + Cl–
Eo = 0.197 V
Fig. Schematic of a saturated
Ag/AgCl electrode
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 Graphical representation of the standard reference potentials
Fig. 17.5 Orbital energy level
(left) and potential (right).
 = – IE (Koopmans)
Fig. 17.6 Orbital energy level for the
SHE (left) and the “absolute” (vacuum)
potential (right).
Potential energy scales plotted such that  = – eE.
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 “Absolute” potential scale for the SHE, EH
779
Established by estimating the energy required to remove electron
from Pt|H2|H+ under standard conditions and transfer it to a vacuum
at rest.  IE = 4.5 eV of a hydrogen atom at standard conditions
H = –IEH = –4.5 eV  EH = H /(–e) = 4.5 V
Ex.
(E vs. SHE) = (E vs. SCE) + 0.242 V
Fig. 17.7 Relationships among SHE, SCE,
and absolute (vacuum) potential scales.
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17.3 MOLECULAR INTERPRETATION OF
ELECTROCHEMICAL PROCESSES
 Energy levels of metals
~ Continuous band of levels
~ Only half of the orbitals are occupied for univalent
metals that contribute one electron per atom
 Energy of the Fermi level, F
~ Highest Occupied Molecular Orbital (HOMO)
F = –  ,
: work function of the metal
cf. Lowest Unoccupied Molecular Orbital (LUMO)
 Potential scale:  = – eE
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 Energy level diagram for the Cu|Cu2+||Ag+|Ag galvanic cell
Fig. 17.8 Electrons from the HOMO in Cu transfer spontaneously
to unoccupied orbitals of Ag+ ion that lie at lower energies or more
positive potentials.
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 Energy level diagram for the Cu|Cu2+||Ag+|Ag electrolytic cell
Fig. 17.9 Electrode potentials and redox reactions for the electrolytic cell
shown in Fig. 17.3. The power supply withdraws electrons from the Ag
electrode and transfers them to empty copper orbitals at higher energies
(more negative potentials). They reduce Cu2+ to Cu at the cathode when
the potential becomes more negative than 0.34 V and Ag is oxidized to
Ag+ at the anode.
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17.4 CONCENTRATION EFFECTS AND THE
NERNST EQUATION
o
Using G = –nFEcell and , G o   nFEcell
o
 nFEcell   nFEcell
 RT ln Q
G = Go + RT ln Q (p. 647) 
 Nernst Equation
o
Ecell  Ecell

RT
ln Q
nF
o
Ecell  Ecell

or
0.0592
log10 Q (at 25oC)
n
n : number of moles of electrons transferred in the overall reaction as written.
 Nernst equation for half-cell reactions (written as reductions)
Ehc  Ehco 
Ex.
0.0592 V
log10 Qhc
nhc
Zn2+ + 2 e–  Zn
nhc = 2, Qhc = 1/[Zn2+]
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 Measuring Equilibrium Constants
o
From G o   nFEcell
and G o   RT ln K
ln K 
nF o
Ecell
RT
or
log10 K 
n
o
Ecell
0.0592 V
(at 25o C)
 pH Meters
(1) Simple pH meter using SHE
Pt | H2(1 atm) | H3O+(var) || H3O+(1 M) | H2(1 atm) | Pt
(anode)
H2(1 atm) + 2 H2O(l)  2 H3O+(var) + 2 e–
(cathode)
2 H3O+(1 M) + 2 e–  H2(1 atm) + 2 H2O(l)
n = 2, Q = [H3
O+(var)]2,
o
cell
and E
SHE
 0 . ( E ’s are the same)
o
hc
Ecell  (0.0592 V / 2) log10 [H 3O  ]2  (0.0592 V) pH
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(2) Smaller and more potable pH meter using two reference electrodes
Ag|AgCl|Cl–(1.0 M) + H3O+(1.0 M)|glass|H3O+(var)||Cl–(sat)|Hg2Cl2(s)|Hg|Pt
Glass indicator electrode || Saturated Calomel electrode (SCE)
Cathode: Hg2Cl2(s) + 2 e–  2 Hg(l) + 2 Cl–(sat)
E1o = +0.2682 V
Anode: 2 Ag(s) + 2 Cl–(1.0 M)  2 AgCl(s) + 2 e– –E2o = –0.2223 V
Dilution: H3O+(1.0 M)  H3O+(var)
Edil = (0.0592 V) pH
Overall: Hg2Cl2(s) + 2 Ag(s) + 2 Cl–(1.0 M) + H3O+(1.0 M)
 2 Hg(l) + 2 AgCl(s) + 2 Cl–(sat) + H3O+(var)
Eo = E1o – E2o
E = Eo– (0.0592 V /2) log [Cl–(sat)]2 + Edil = Eref + (0.0592 V) pH
pH 
E  Eref
0.0592 V
where Eref = Eo– (0.0592 V) log [Cl–(sat)]
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 Potential for the saturated calomel electrode:
E(SCE) = E1o – (0.0592 V) log [Cl–(sat)] = + 0.241 V
Fig. 17.10 Schematics
of an early pH-meter.
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Fig. 17.11 A combination pH probe
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17.5 MOLECULAR ELECTROCHEMISTRY
 Electrochemical Organic Synthesis
 Monsanto process for the production of adiponitrile
CH2=CHCN + 2 e–  2(CH2=CHCN)– + 2 H+  NC(CH2)4CN
Adiponitrile  hexamethylenediamine  nylon
Electrochemical reactors with steel electrodes coated with Cd film
No need to use powerful but toxic redox agents
Suitable for synthesis of pharmaceuticals (speed, selectivity)
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 Enzyme-Based Electrochemical Sensors
 Monitoring blood sugar levels for diabetics
(1) Glucose sensor test strip monitors…
Oxidation of glucose by glucose dehydrogenase, GDH(FAD)
(2) Implanted device monitors…
Oxidation of glucose by glucose oxidase, GOx(FADH2)
FAD, FADH2: reduced and oxidized forms of cofactors
 FAD (Flavin Adenine Dinucletide): redox cofactor
 Cofactor: non-protein chemical compound that is required
for the protein's biological activity.
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 Glucose sensor test strip contains…
① A capillary drawing in the blood sample
② A thin film carbon working electrode
③ A thin film Ag/AgCl electrode
(counter-reference electrodes)
④ Two planar auxiliary electrodes
(fill-detection electrodes)
Fig. 17.12 Glucose sensor test strip
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(1) Glucose sensor test strip monitors the concentration of
glucose by measuring the total charge transferred to the
carbon electrode that results from the oxidation of glucose
by GDH, which converts an alcohol group into a ketone:
 glucononlactone + 2 H++ 2 e–
② glucose 
GDH
(2) Reduction half-reaction occurs at the Ag/AgCl reference electrode:
 2 Ag
③ 2 Ag+ + 2 e– 
(3) Overall set of reaction:
GDH
 glucononlactone + 2 Ag + 2 H+
glucose + 2 Ag+ 
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 Mediators:
~ Molecules facilitating electron transfer over large
distances by a series of efficient electron transfer reactions
~ Mediators used in this particular sensor are coordination
complexes of osmium (Os2+/Os3+)
~ Redox potential is adjusted so that its oxidized form could
be reduced by the enzyme, but not at the Ag/AgCl electrode.
Fig. 17.13 Coupled redox reactions in the glucose test strip.
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 Implantable glucose sensors
– Continuous measurement of glucose
– Transmitting signals of concentration change to a monitor
– ‘Wired’ enzyme:
Fe2+/Fe3+ (or Os-based) redox couple covalently attached
to the polypeptide chain of glucose oxidase (GOx)
– GOx is incorporated into the mediator-bearing polymer
 Electrogenerated Chemiluminescence (ECL)
– Analytical technique for detecting biologically important molecules
– Emission from excited states molecular electronic states produced
by charge transfer reactions between radical anions and radical
HOMO cations:
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A  e  
 A 
LUMO
HOMO
LUMO
HOMO
 D   e 
D 
 A*  D
A   D  
 A  h
A* 
Fig. 17.14 Generation of a radical cation
and a radical anion in an ECL experiment.
A  (or D ) is formed when the electrode
potential is made more negative (or positive)
than the LUMO (or HOMO) of A (or D).
Fig. 17.15 Reaction of R– and R+ to form R*, which emits light,
in an ECL experiment.
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 Cyclic voltammetry (CV)
~ Measures the current that flows as the electrode potential
is scanned back and forth
~ Identify redox active materials
~ Locate the potentials at which various oxidations and
reduction reactions occur
~ Information about the kinetics of electrochemical reactions
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Reduction of PM to PM
Oxidation of PM toPM


Fig. 17.16 Cyclic voltammogram for an ECL dye, PM 567.
Cathode: Redox potential for the PM| couple: –1.4 V LUMO
Anode: Redox potential for the |PM couple: 1.0 V HOMO
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 Calculating the maximum amount of energy available to
create an excited state, and the energy of the emitted
photon from the reaction
~ Energy available between two redox couples:
  eE  e  Ecathode  Eanode   e( 1.4  1.0)  2.4 eV
 Absorption peak at 514 nm  2.4 eV
~ Energy available from the recom
bination reaction is sufficient to
create the excited state
 Fluorescence peak at 567 nm
~ Typical shift to longer wavelength
in aqueous solution
Fig.
17.17. Chemistry
Absorption and
General
II fluorescence spectra of PM 567.
793
]3–
]4–]
Cyclic voltammogram for [Fe(CN)6 /[Fe(CN)6
redox couple measure with respect to a Ag/AgCl reference electrode.
The standard reduction potential for this system at pH=7 is 0.43 V.
EXAMPLE 17.10
Reduction
[Fe(CN)6]4–]
Re-oxidation
[Fe(CN)6]3–]
of
of
[Fe(CN)6]3–
[Fe(CN)6]4–
to
to
The redox potential is determined graphically by locating the
midpoint between the two peaks, in this case 0.43 V.
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 ECL reaction in aqueous solution
Tris(bipyridine)ruthenium(II) cation: Ru(bpy)32+
Coreactant: Tripropylamine, N(CH2CH2CH3)3 (or Pr2NCH2CH2CH3)
Pr2NCH2CH2CH3
 Overall sequence of the reaction:
Ru(bpy)32 
 Ru(bpy)33  e 
Ru(bpy)32+
Pr2 N(CH 2 CH 2 CH 3 ) 
 Pr2 N(CH 2 CH 2 CH 3 )   e 
Pr2 N(CH 2 CH 2 CH 3 )  
 Pr2 N( CHCH 2 CH 3 )  H 
  Ru(bpy)32  *  Pr2 N(  CH 2 CH 2 CH 3 )
Ru(bpy)33  Pr2 N(CHCH 2 CH 3 ) 
 Ru(bpy)  *
2
3

 Ru(bpy)32  h
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 Applications of ECL
 Ru(bpy)32+ is attached to the target molecules
Antigens, DNA, RNA
 Antibody-based assays:
Measure insulin levels
Measure estrogen and testosterone levels
Detect the hepatitis and HIV viruses
 Advantages of ECL
Stable labels,
No background emission
Very sensitive (10–12 M)
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EXAMPLE 17.15 An ECL molecule emits light at 600 nm and has a potential
associated with the LUMO located at –1.05 V. The energy of this LUMO
is 1.05 eV. Calculate the redox potential associated with the HOMO?
 = h = hc / 
= (6.62610–34 J s)(3108 m s–1) / (60010–9 m)
 = – eE(V)
= 3.310–19 J = 2.06 eV
= LUMO – HOMO = 1.05 eV – HOMO
 HOMO = – 1.01 eV
~ with an associated potential
located at +1.01 V
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 Chromophore
A molecule (or a semiconductor) that absorbs light
~ Lower energy level: HOMO (Valence Band, VB)
~ Higher energy band: LUMO (Conduction Band, CB)
Fig. 17.18 Photoexcitation promotes
electrons to higher energies, which
make them stronger reducing agent,
capable of reducing an acceptor A to
A–. The vacancies (holes) left behind
can now accept electrons, which make
them stronger oxidizing agents,
capable of oxidizing a donor D to D+.
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 Direct photoelectrochemical water splitting
by wide bandgap semiconductors (TiO2 or SrTiO3)
Fig. 17.19 The potential of the conduction band of wide bandgap
semiconductors is sufficiently negative to reduce hydrogen ions and
that of the valence band sufficiently positive to oxidize water.
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 Overall process with CB electrons reducing protons and VB holes
oxidizing water:
4 H+(aq) + 4 e–  2 H2(g)
2 H2O(l)  O2(g) + 4 H+(aq) + 4 e–
 Standard Gibbs free energy of formation of water:
H 2 (g )  12 O2 (g ) 
 H 2 O(l )
Gfo  237.18 kJ mol1
 The maximum amount of energy available becomes the difference
in energies of the hydrogen and oxygen redox levels:
G o  nFE o  (1 mol)(96,585 kJ mol1eV 1 )(1.299 eV)  125.3 kJ
per mole of electrons.
 Water splitting is a four-electron redox process that proceeds through
a series of electron transfer reactions involving several intermediates.
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Fig. 17.20 Direct photoelectrochemical water splitting by colloidal
TiO2, shown with hydrogen and
oxygen evolution catalyst particles
attached.
 Wide bandgap semiconductors (TiO2, SrTiO3)
– Absorb radiation only in the UV region
 Only 10% of the energy provided by sunlight
– C B with potentials sufficiently negative to reduce water
– VB located at much more positive potentials than necessary
to oxidize water  Most of energy absorbed is lost as heat
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CONNECTION TO ENERGY
798
 Solar Energy Conversion
 Artificial Photosynthesis
Fig. Direct photoelectrochemical
water splitting system at visible
region with the addition of a dye
sensitizer and a mediator
~ LUMO energy > CB energy for injection of electrons (reduction)
~ Redox potential of HOMO must be sufficiently positive to oxidize water.
~ Redox mediators (Small Pt and RuO2 particles) work as catalysts
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 Dye-Sensitized Solar Cells (DSC)
 Michael Grätzel (Swiss, 1944 - )
~ Discovered a new type of dye-sensitized solar cell:
Semiconductor (TiO2) + Sensitizer (Ru(bpy)3) + Mediator (I–/I3– couple)
General Chemistry II
Fig. Schematic of a Grätzel cell
799
 Sensitizers:
~ Derivatives of Ru(bpy)3 in which one of more of bipyrimidine
ligands are replaced by isothocyanate (–N=C=S) ligands
 Extends the absorption onset to longer wavelengths
~ Conversion efficiency: incident photons to electrons
~ Solar spectrum: represented by the spectrum of a blackbody
at the temperature of sun
Fig. Absorption of radiation
By the sensitizers, RuL2(NCS)2
and RuLL’(NCS)3
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 Redox mediator
~ Transports electrons from the cathode to the radical
cation created by photoexcitation and charge injection
~ Regenerating the ground state of the sensitizer
~ Preventing direct electron-hole pair recombination.
h
S 
 S  e  (TiOcb
2 )


I  e 
 I  I 2 
I

3
S  I3 
 S  I2  I
charge injection
mediator reduction (I  /I3 redox couple)
sensitizer regeneration
General Chemistry II
17
ELECTROCHEMISTRY
CHAPTER
17.6 Batteries and Fuel Cells
17.7 Corrosion and Corrosion Prevention
17.8 Electrometallurgy
17.9 Electrolysis of Water and Aqueous Solutions
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17.6 BATTERIES AND FUEL CELLS
 Batteries
~ Primary cells, Secondary cells (rechargeable)
 Zinc-carbon “Dry” cell
(Leclanché cell): 1.55 V
Cathode: graphite rod
MnO2 + graphite powder
(large surface area)
2 MnO2(s) + 2 NH4+(aq) + 2 e–
 Mn2O3(s) + 2 NH3(aq) + H2O(l)
Salt bridge:
moist paste of ZnCl2 + NH4Cl
Anode: zinc shell
Zn(s)  Zn2+(aq) + 2 e–
Fig. 17.21 Leclanché “dry” cell
Overall reaction:
Zn(s) + 2Chemistry
MnO2(s) + II2 NH4+(aq)  Zn2+(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l)
General
802
 Disadvantage
[NH4+] decreases with time  Battery’s voltage falls with time
Zinc anode corrodes as it oxidizes
Electrolyte leaks out
 Alkaline dry cell:
KOH instead of NH4Cl, 1.5 V (Eo = 1.364 V)
Cathode:
2 MnO2(s) + H2O(l) + 2 e–  Mn2O3(s) + 2 OH–(aq)
Anode:
Zn(s) + 2 OH–(aq)  Zn(OH)2(s) + 2 e–
Overall:
Zn(s) + 2 MnO2(s) + H2O(l)  Zn(OH) 2(s) + Mn2O3(s)
No dissolved species: Constant concentration  steadier voltage
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 Zinc-mercuric oxide cell
Very stable, 1.34 V, Camera, watch, calculators
Electrolyte: 45% KOH solution
Cathode: Steel & HgO(s)
HgO(s) + H2O(l) + 2 e–
 Hg(l) + 2 OH–(aq)
Anode: Mixture of Hg & Zn
Zn(s) + 2 OH–(aq)
 Zn(OH)2(s) + 2 e–
Overall:
Zn(s) + HgO(s) + H2O(l)
 Zn(OH)2(s) + Hg(l)
Fig. 17.22 A zinc-mercury oxide dry cell.
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802
 Lithium battery
Replacing zinc-mercuric oxide dry cell due to environmental reason
Battery voltage: 3 V
Electrolyte: LiClO4 dissolved in a mixture of propylene carbonate
and 1,2-dimethoxyethane
(anode)
Li  Li+ + e–
(cathode) (Mn4+)O2 + Li+ + e–  (Mn3+)O2(Li+)
(overall)
Li + (Mn4+)O2  (Mn3+)O2(Li+)
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 Rechargeable Batteries
 Characterized by:
Voltage, maximum current available, total energy stored
 Capacity: Amount of charge [current (i)  time (t)] available
in a battery before it must be recharged
 100 watt hour (Wh) battery: delivers 5 A for 20 hours
 3500 mAhr battery: provides 3.5 A for an hour or 1 A for 3.5 hour
Energy available from a battery before it must be recharged
= capacity  voltage = itV = (3500 mAhr)(10 V) = 35 J
 Volumetric energy density = energy per unit volume, Wh/L
Gravimetric energy density = energy per unit weight, Wh/kg
Ex. Modern Li-ion laptop batteries achieve 200 Wh/kg
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Fig. 17.23 Ranges of typical energy densities for different classes of batteries.
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 Lead-acid storage battery:
6  2.0 V = 12 V
Cathode: PbO2(s) + SO42–(aq) + 4 H3O+(aq) + 2 e–  PbSO4(s) + 6 H2O(l)
Anode:
Pb(s) + SO42–(aq)  PbSO4(s) + 2 e–
--------------------------------------------------------------------------------------------------------Overall: Pb(s) + PbO2(s) + 2 SO42–(aq) + 4 H3O+(aq)  2 PbSO4(s) + 6 H2O(l)
0
+4
common electrolyte
+2
Many cycles of charge-discharge, Large current as large as 100 A,
Heavy (low energy density)  limits the range of the electric vehicles
Fig. 17.24 A lead-acid storage battery
consists of several cells connected in series.
The electrodes are both constructed from
lead grids filled with spongy Pb (cathode)
and PbO2 (anode) and the electrolyte is
sulfuric acid.
General Chemistry II
 Lithium-ion battery
804
High energy density (light Li)
Much safer than lithium battery (no elemental Li, only Li+ ions)
Electrodes: LiCoO2 (cathode), LiyC6 (graphite anode)
Single Li+ ions shuttle back and forth:
LiC6  Li+ + e– + C6
(anode)
LiMn+O2 + Li+ + e–  Li2M(n–1)+O2
(cathode)
------------------------------------------------------------n+
(n–1)+O + C
2
6
General ChemistryLiC
II 6 + LiM O2  Li2M
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 Alkali metal batteries (Na/S battery)
High energy density, high T (250°C) cells
Cathode: S + nS + 2 e–  Sn+12–
Anode:
2 Na  2 Na+ + 2 e–
------------------------------------------------------Overall: S + nS + 2 Na  Sn+12– + 2 Na+
-alumina (NaAl11O17):
Ion-conducting solid electrolyte
Only Na+ can migrate
CaF2 solid electrolyte:
Only F– can migrate through
Fig. 17.26 Schematic of a sodium-sulfur battery.
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805
 BOB (Big Old Battery): Na/S battery of the size of a house
~ Supplies 4 MW of power for 8 hrs overnight in remote (60 miles)
small village, Presidio, Texas (population, ~ 5000)
~ $25 million project, built by NGK Insulators of Japan
Fig. Na/S battery house built in Presidio, Texas.
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 Fuel Cells
 Batteries:
Closed system, recharged or discarded after use
 Fuel cell: Energy converter
~ Continuous operation,
~ Supplying reactants and removing products
 Two classes of fuel cells
(1) Polymer electrolyte membrane (PEM) fuel cells
(2) Solid oxide fuel cells (SOFC)
General Chemistry II
806
 Polymer electrolyte membrane (PEM) fuel cell
Anode:
H2(g) + 2 H2O(l)  2 H3O+ + 2 e– Eo = 0.000 V
Cathode: O2(g) + 4 H3O+ + 4 e–  6 H2O(l)
Eo = 1.229 V
--------------------------------------------------------------------------------------Overall:
2 H2(g) + O2(g)  2 H2O(l)
Ecello = 1.229 V
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 Maximum work available from combustion:
807
 wmax   qP  (1  Tl / Th )H fo
 : Carnot efficiency of an internal combustion engine (~45%)
qP: heat transferred at constant pressure
Tl: temperature of the exhaust gas
Th: operating temperature of the internal combustion engine

 wmax  ( 0.4)( 241.8 kJ mol 1 )  96 kJ mol 1
 Maximum work available from a fuel cell:
 wmax  Gfo  228.6 kJ mol1
 Comparison of the “Tank-to-wheel” efficiency
 Mechanical efficiency of the engines: combustion(75%), fuel cell(90%)
Overall efficiency fuel cell vehicle
(0.9)(228.6 kJ mol1 )

 2.89
Oveall efficiency internal combustion engine vehicle
(0.75)(96 kJ mol1 )
General Chemistry II
 Solid oxide fuel cell (SOFC)
808
~ Large scale stationary power generation for hospitals
and remote towns, providing 250 kW~1 MW power
Electrolyes: Oxide ceramics, ZrO2, doped with Y2O3
H 2 (g ) 
1
2
O2 (g ) 
 H 2 O(l )
H fo  245.8 kJ mol1
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17.7 CORROSION AND CORROSION PREVENTION
Spontaneous electrochemical reactions of iron
~ “short-circuited” galvanic cell
 In the absence of oxygen; slow and not serious corrosion.
Cathode:
2 H2O(l) + 2 e–  H2(g) + 2 OH–(aq)
Anode:
Fe(s)  Fe2+(aq) + 2 e–
------------------------------------------------------------------------------------Overall:
Fe(s) + 2 H2O(l)  Fe2+(aq) + H2(g) + 2 OH–(aq)
 With both oxygen and water Fig 17.30
Cathode:
1/2 O2(g) + 2 H3O+(aq) + 2 e–  3 H2O(l)
Anode:
Fe(s)  Fe2+(aq) + 2 e–
------------------------------------------------------------------------------------Overall: Fe(s) + 1/2 O2(g) + 2 H3O+(aq)  Fe2+(aq) + 3 H2O(l)
General Chemistry II
809
Further oxidization of Fe2+(aq) migrated to cathode to form rust
2 Fe2+(aq) + 1/2 O2(g) + (6 + x) H2O(l)  Fe2O3·xH2O(s) +4 H3O+(aq)
rust
Overall corrosion reaction:
2 Fe(s )  3/2 O 2 (g )  x H 2 O(l ) 
 Fe 2 O 3  xH 2 O(s )
Fig. 17.30 Schematic of iron corrosion. Hydrogen ions move through
the hydrated rust pile.
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– Pitting on the anode (oxygen-poor areas under the paint)
– Rust on the cathode (oxygen-rich exposed area)
– Dissolved salts increase the conductivity.
– High acidity due to air pollution or CO2 speeds up the
reduction at the cathode.
 Passivation: Protective thin metal oxide layer
Spontaneous aluminum oxidation (Al2O3)
Stainless steel (alloy of iron and chromium)
Rust preventing (superficial oxidation) paints containing
K2Cr2O7 and Pb3O4
General Chemistry II
810
 Sacrificial Anode: A piece of Mg in contact with Fe
Fe2+(aq) + 2 e–  Fe(s)
Eo = –0.41 V
Mg2+(aq) + 2 e–  Mg(s)
Eo = –2.39 V
~ Mg2+ is much harder to reduce than Fe2+
Passivation by molten zinc
Sacrificial Mg anode
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17.8 ELECTROMETALLURGY
 Extractive metallurgy
~ Over the Stone Age ~ 6000 yrs. ago (Near East)
~ Elemental form ~ Au, Ag, Cu, Fe in meteorites
o
~ Metal oxides ~ Gf  0
~ Moderate heating of HgS or HgO ( Gfo  29 kJ mol1 )
 Pyrometallurgy
Reduction of metal oxides in furnace heated by coke at high temp
Smelting ~ chemical change + melting
CuO, PbO, NiO, Fe2O3, SnO2, ZnO
~ Driving force for the overall reaction:
C(s) + O2(g)  CO2(g)
G o  394 kJ mol1
~ Al, alkali metals ~ high free energy w.r.t. their oxides (ores)
General Chemistry II
811
 Aluminum
3rd most abundant element in earth’s crust
~ 8.2 % by mass (1st: O, 2nd : Si)
Bauxite ~ hydrated aluminum oxide containing 50-60 % Al2O3,
1-20 % Fe2O3, 1-10 % silica, traces of Ti, Zr, V,.. oxides,
20-30 % water
 Properties of Al
– Low density ~ vehicles
– High conductivity ~ aluminum wires
– Resistance to corrosion ~ chassis
– Gets stronger at subzero temp
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 Bayer process (purification of bauxite)
Amphoteric oxides soluble in strong bases:
Al2O3(s) + 2 OH–(aq) + 3 H2O(l)  2 Al(OH)4–(aq)
Iron(III) oxide is not soluble in strong bases
~ Separated by filtration
Cooling down to supersaturation:
2 Al(OH) 4– (aq)  Al2O3·3H2O(s) + 2 OH–(aq)
Hydrated water removed by calcining at high temp (1200oC)
 Sir Humphrey Davy (1809)
~ Discovered as an alloy of iron and proved its metallic property
 H. C. Örsted (1825) ~ Obtained relatively pure form
AlCl3(s) + 3 K(Hg)x(l) 3 KCl(s) + Al(Hg)3x(l)
~ Hg removed by distillation
 Charles Hall & Paul Héroult (1886) ~ Invented modern method
General Chemistry II
812
 Hall-Héroult process
 Cathodic deposition of Al by electrolysis of molten cryolite (Na3AlF6)
containing dissolved Al2O3(m.p. 2050 oC)
(Lowering m.p. of cryolite from 1000 oC  950 oC)
 50,000-100,000 A current
 100 electrolysis cells connected in series
Anode ~ graphite
Cathode ~ rectangular steel box (6 m 2 m 1 m)
Anode: C / Al2O3 in Na3AlF6(l) / Al(l) / steel box : Cathode
Overall cell reaction: 2 Al2O3 + 3 C  4 Al + 3 CO2
Al(l) denser than the melt  collected at the bottom
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Fig. 17.32 An electrolytic cell used in the Hall-Heroult
process for the commercial production of aluminum.
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813
 Magnesium
 Sources of Mg
Dolomite (CaMg(CO3)2), Carnallite (KCl MgCl2 6H2O),
2nd most abundant cation in the seawater
 Separation of Mg2+ from the seawater
(1) Calcined dolomite
Addition of low-cost base (calcined dolomite) to seawater
CaMg(CO3)2(s)
High T


CaO MgO(s) + 2 CO2(g)
CaO MgO(s) + Mg2+(aq) + 2 H2O(l)  2 Mg(OH)2(s) + Ca2+(aq)
Mg(OH)2 is least soluble compared to hydroxides of Na+, Ca2+, K+
 K sp (Ca(OH)2 )  5.5  106 , Ksp (Mg(OH)2 )  1.2  1011 
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(2) Oyster shells (Texas seashore)
 calcined to CaO  added to seawater to produce Mg(OH)2
 Magnesium carbonate
(Coating table salt to prevent caking. Antacid remedies)
Mg(OH)2 + CO2  MgCO3 + H2O
 Magnesium chloride and reduction to pure magnesium
Mg(OH)2(s) + 2 HCl(aq)  MgCl2(aq) + 2 H2O(l)
Electrolysis of molten MgCl2 (m.p. 708 oC) in a large steel cell
MgCl2(l)  Mg(l) + Cl2(g)
cathode: steel, anode: graphite
Cl2(g) recycled to produce HCl
 Magnesium metal
Flashbulbs (till 1918), Sacrificial anode, Reducing agent
General Chemistry II
814
Fig. 17.33 The production of Mg(OH)2 starts with the addition of lime (CaO) to
seawater. Reaction of Mg(OH)2 with MgCl2, which, after drying, is reduced to
produce Mg metal.
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 Electrorefining
814
~ Impure copper slabs (anode) alternating with thin sheets
of pure copper (cathode) in CuSO4/H2SO4(aq)
~ Impurites
Nickel  dissolves into solution
Silver, Gold  fall down as metallic slimes
Fig. 17.34 Electrolytic refining of copper.
 Electroplating
~ Pure chrome (anode), Metal to be plated (cathode)
in hot bath of H2SO4 / H2CrO4
–5
–2
~
Thickness
General Chemistry: II2.5 x 10 ~ 10 cm
816
17.9 ELECTROLYSIS OF WATER AND AQUEOUS
SOLUTIONS
 Electrolysis of pure water ~ Inert electrodes (Pt)
Eo = 0.00 V
Cathode: 2 H3O+(aq) + 2 e–  H2(g) + 2 H2O(l)
+
–
Anode: 3 H2O(l)  (1/2) O2(g) + 2 H3O (aq) + 2 e – Eo = –(+1.229 V)
----------------------------------------------------------------------------------------------Overall: H2O(l)  H2(g) + (1/2) O2(g)
Eo = –1.229 V < 0
Since [H3O+] = [OH–] = 1.0 10–7 M, P(H2) = P(O2) = 1 atm at 25°C,
E(cathode)  Eo (cathode) and E(anode)  Eo (anode)
E(cathode) = Eo (cathode) – (0.0592 V / nhc) log Qhc
= 0.00 – (0.0592 V / 2) log {P(H2)/[H3O+]2}
= 0.00 – (0.0592 V / 2) log {1/ (10–7)2}
= – 0.414 V = E (H3O+(10–7 M)|H2(1 atm))
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E(anode) = Eo (anode) – (0.0592 V / nhc) log Qhc
= 1.229 – (0.0592 V / 2) log {P(O2)/[H3O+]2}
= 1.229 – (0.0592 V / 2) log {1/ (10–7)2}
= 0.815 V = E(O2(1 atm),H3O+(10–7 M)|H2O)
E = E(cathode) – E(anode) = – 0.414 – 0.815 V = –1.229 V ( = Eo )
Since E < 0, G > 0  nonspontaneous
 needs an external voltage, Decomposition potential
(for water : 1.229 V)
In reality, the voltage drop between the two electrodes
and the effect of overvoltage need to be considered.
General Chemistry II
817
 Electrolysis of an electrolyte solution, 0.1 M
NaCl(aq)
Possible half-cell reactions at electrode:
Cathode: Na+(0.1 M) + e–  Na(s)
2 H3
O+(10–7
M) + 2
e–
 H2(g) + 2 H2O(l)
Eo (Na+|Na) = –2.71 V
Eo(H3O+|H2) = 0.00 V
2 Cl–(0.1 M)  Cl2(g) + 2 e–
– Eo(Cl2|Cl–) = –1.36 V
+
–7
–
6 H2O(l)  O2(g) + 4 H3O (10 M) + 4 e – Eo(O2,H3O+| H2O) = –1.229 V
Anode:
[H3O+] = [OH–] = 1.0 10–7 M, P(H2) = P(Cl2) = P(O2) = 1 atm at 25°C
 Reduction potential for the first reaction:
E(Na+|Na) = Eo (Na+|Na) – (0.0592 V / nhc) log Qhc
= –2.71 – (0.0592 V / 1) log {1/[Na+]}
= –2.71 – (0.0592 V / 1) log {1/(0.1)}
= – 2.71 – 0.06 = – 2.77 V
Smaller than – 0.414 V = E[H3O+(10–7 M)|H2(1 atm)]
 Reduction of Na+(aq) impossible !
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 Reduction potential for the third reaction:
E(Cl2|Cl–) = Eo (Cl2|Cl–) – (0.0592 V / nhc) log Qhc
= 1.36 – (0.0592 V / 1) log {[Cl–]/P(Cl2)}
= 1.36 – (0.0592 V / 1) log {(0.1) / 1} = 1.42 V
Larger than 0.815 V = E(O2(1 atm),H3O+(10–7 M)|H2O)
 reduction of Cl2(g) possible !
 Therefore, the actual half-reactions are:
Cathode:
2 H3O+(10–7 M) + 2 e–  H2(g) + 2 H2O(l)
Anode:
6 H2O(l)  O2(g) + 4 H3O+(10–7 M) + 4 e–
General Chemistry II
817
 Replace 0.1 M NaCl solution with 0.10 M NaI solution
Na+ ions still will not be reduced.
E(I2|I–) = Eo (I2|I–) – (0.0592 V / nhc) log Qhc
= 0.535 – (0.0592 V / 1) log [I–]
= 0.535 + 0.059 = 0.594 V
More negative than 0.815 V = E(O2(1 atm),H3O+(10–7 M)|H2O)
 Oxidation of 0.10 M I– occurs in preference to the oxidation of water
Therefore, the actual half-reactions are:
Cathode:
2 H3O+(10–7 M) + 2 e–  H2(g) + 2 H2O(l)
Anode:
2 I– (0.10 M)  I2(s) + 2 e–
Overall:
2 H3O+(10–7 M) + 2 I– (0.10 M)  H2(g) + I2(s) + 2 H2O(l)
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The intrinsic cell voltage is
E = E(cathode) – E(anode) = – 0.414 – 0.594 V = –1.008 V < 0
(1) H2(g) and I2(s) will be generated by applying a potential greater than
the decomposition potential of the solution, which is 1.008 V.
(2) Concentration of I– begin to decrease as the electrolysis proceeds,
making the potential of the (I2|I–) couple more positive.
(3) When the concentration of I– reaches about 210–5 M, E(I2|I–) = 0.815 V,
the external voltage required to maintain electrolysis would have to be
increased to 1.229 V.
(4) At this point, water will start to be electrolyzed, and oxygen will be
produced at the anode.
Electrolysis of neutral aqueous solutions
1. A species can be reduced only if E(reduction) > –0.414 V
2. A species can be oxidized only if E(reduction) < 0.815 V
General Chemistry II
10 Problem Sets
For Chapter 17,
6, 14, 34, 46, 54, 60, 70, 76, 88, 104
General Chemistry II
42

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