Electrochemistry Lecture 1 Introduction and Overview of


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CHE 729 Electrochemistry
Lecture 1 Introduction and Overview of
Electrode Processes
Dr. Wujian Miao 1
Electrochemistry
The study of chemical reactions that involve electron or charge transfer (e.g., ions).
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e
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Electroanalytical Chemistry
A large group of analytical methods (Quantitative and Qualitative) based on electrochemistry. For example:
Voltammetry, Potentiometry, Electrogenerated Chemiluminescence (ECL), Electrochemical Scanning Microscopy (SECM), Electrochemical Biosensors.
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Some Advantages of Electroanalytical Methods • fast • inexpensive • in situ (pH, glucose) • information about oxidation states • stoichiometry • rates of diffusion, electron transfer, • equilibrium constants 5

Energy Conversions with Cells

Chemical potential energy

galvanic cell
Electrical potential energy
electrolytic cell

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(a) A galvanic electrochemical cell at open circuit; (b) a galvanic cell doing work; (c) an electrolytic cell.
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Redox Processes - electron transfer
Cu(s)  2e Oxidation (-e) [Galvanic Cell] Cu2 (aq) Reduction (+e) [Electrolyte Cell]
2Ag+ (aq)  2e Reduction (+e) [Galvanic Cell] 2 Ag (s) Oxidation (-e) [Electrolyte Cell]
Cu(s)  2Ag+ (aq) GalvanicCell  Cu2 (aq)  2 Ag(s) Electrolytic Cell
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Anode vs Cathode
Anode: electrode with oxidation process Cathode: electrode with reduction process
• Depends on the type of electrochemical cell • Determined by the redox process 9
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Movement of electrons and ions in a cell—Functions of salt bridge
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Liquid Junction the interface between electrochemical solution and the salt bridge (liquidliquid interface)  could develop a small junction potential
Anode (Negtive electrode): Cd  Cd2  2e Cathode (Positive electrode): AgCl  e  Ag  Cl
Cell without a salt bridge 11
Origin of Liquid Junction Potentials
• HCl (1.0 M)|HCl (0.01 M) Differences in ion mobility give rise to junction potentials. Unequal distribution. Solution: salt bridge containing a high concentration of KCl or NaNO3
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Membrane

Shorthand Cell Notation
Convention
• Anode—on the left • Cathode-right • |--interface • ||--salt bridge (two
liquid-liquid interfaces • Activity/concentration/
pressure included
Zn|ZnSO4 (aZn2+ = 0.0100 M)||CuSO4(aCu2+ = 0.0100M)|Cu (Anode, negative) (salt bridge) (cathode, positive)
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(Anode, -) Cd|CdCl2 (0.0167M)|AgCl|Ag (Cathode, +) Pt, H2 (p = 1atm)|H+ (0.01 M), Cl- (0.01 M),AgCl(sat'd)|Ag
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Electrode Potentials—Nernst Equation

Ox  ne  Red ("half cell" reaction)

E  E0  RT ln aOX nF aRed

{Lnb  logb  log1b0 } e log1e0

 E0  8.316298.15 log aOX  E0  0.05918 log aOX

n96485.31log 2.718 aRed

n

aRed

where E0 is the standard potential (i.e, when aOX = aRed  1)

R is the gas constant (8.316 J mol-1K-1), T is the temperature (K),

n is the number of electrons transferred, F Faraday constant 96485.31 C mol-1

aX   X[X]  X is the activity coefficient of solute X.  X varies with the presence of other ions (ionic strength

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Formal/Conditional Potentials

E  E0  RT ln aOX  E0  RT ln  OX[OX]

nF aRed

nF  Red[Red]

 {E0  RT ln  OX }  RT ln [OX] nF  Red nF [Red]

 E0'  RT ln [OX]  E0'  RT ln [Red]

nF [Red]

nF [OX]

(E0'  formal/conditional potential)

e.g., Hydrogen 2H + 2e = H2 (g) E0 =0,
E0’ = -0.005 V in 1 M HCl, HClO4

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Examples
Write Nernst expressions for the following half-cell reactions:
(1) Zn2+ + 2e = Zn (s) (2) Fe3+ + e = Fe2+(s) (3) 2H+ + 2e = H2(g) (4) MnO4- + 5e + 8H+ = Mn2+ +4H2O (5) AgCl(s) + e = Ag(s) + Cl-
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(1)E  E0  RT Ln [Zn2+ ]  E0  0.0592 log [Zn2+ ]

2F [Zn]

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(2)E  E0  0.0592 log [Fe3+ ]

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[Fe2+ ]

(3)E  E0  0.0592 log [H+ ]2 2 pH2

(4)E  E0  0.0592 log [MnO4- ][H+ ]8  E0  0.0592 log [MnO4- ][H+ ]8

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[Mn2+ ][H2O]4

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[Mn2+ ]

(5)E  E0  0.0592 log [AgCl(s)]  E0  0.0592 log 1

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[Ag(s)][Cl- ]

1

[Cl- ]

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Electrochemical Cell Potentials —Nernst Equation
• The Nernst equation can be derived from the equation linking free energy changes to the reaction quotient:
G  G0  RT ln Q where Q is the reaction quotient  G  nFE E is the electrochemical cell potential
G0  nFE0 E0 is the standard electrochemical cell potential E  E0  RT ln Q
nF  E0  0.0592 l o g Q at 298 K (25 o C)
n
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E(or Ecell )  Ecathode()  Eanode()

For aA  bB  ...  pP  sS  ...

(n electrons involved)

Q



aPp



a

s S

.

.

.

aAa  aBb ...

When the reaction reaches equalibrium ΔG  0 (E  0), Q  Keq

 G0  RT ln Keq

or -nF E0  RT ln Keq

ln Keq  nFRTE0 or lgKeq  0n.05E902 (at 298K)

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Standard (Normal) Hydrogen Electrode (SHE/NHE)
Can't measure potential on each electrode independently – only differences
2H+ (a =1.00 M)+ 2e = H2 (g, p = 1.00 atm) Pt, H2 (g, p = 1.00 atm) | H+ (a =1.00 M)||... Assigned 0.000 V at any Temperature. Can be anode or cathode. Pt-electron carrier, does not take part in rexn. Cumbersome to operate. Ecell = EM  ESHE
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Practical Reference Electrodes (Secondary Ref. Electrodes)

Aqueous

Non-aqueous

 SCE (Saturated Calomel Electrode)
 Ag/AgCl (KCl, x M)  Ferrocene methanol--
Fc-CH2OH—water soluble,  added directly to the analyte solution.

 Ag/Ag+ (x M AgNO3)  Pseudo (Quasi)
references
Pt, Ag wires
 Ferrocene—Internal standard potential ref.
 added directly to the analyte solution.

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Saturated Calomel Electrodes (SCE)
 Hg(l)|Hg2Cl2|KCl(Sat’d)  E0 = 0.241 V vs. SHE @ 25°C

Hg2Cl2 (s) + 2e = 2Hg(l) + 2Cl- EH0g2Cl2 /Cl- = 0.268 V vs SHE

E

 E0  RT ln 1  0.268  0.0592 lg a

Hg2Cl2 /Cl-

Hg2Cl2 /Cl- 2F (a )2

Cl 

Cl 

(Satured KCl, [Cl- ]~4.5 M)

Electrode Hg(l)/Hg2Cl2(s)/KCl (0.1 M)
Hg/Hg2Cl2(s)/KCl (1 M) Hg(l)/Hg2Cl2(s)/KCl (sat'd) Hg(l)/Hg2Cl2(s)/ NaCl (sat'd)

Acronym
NCE SCE SSCE

Potential vs. SHE
0.3337 0.2801 0.2412 0.2360

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Ag/AgCl (x M KCl)
 Ag wire coated with AgCl(s), immersed in NaCl or KCl solution
 E0 = 0.222 V vs. SHE @ 25°C
AgCl + e = Ag + ClEAg/ AgCl  E0  0.0592 lg a1
Cl 
 E0  0.0592 lg aCl aCl , EAg / AgCl  .
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Sat’d AgCl/KCl

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Electrochemistry Lecture 1 Introduction and Overview of