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Stewart SmithBiosensors and InstrumentationBeijing University of Posts and
Telecommunications 2019
Introduction to Electrochemistry
Lecture 4
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Stewart SmithBiosensors and InstrumentationBeijing University of Posts and
Telecommunications 2019
Summary
• Redox reactions
• Standard electrode potential
• Control of electrode reactions
• 3-Electrode cell and Reference electrodes
• Ion sensitive electrodes
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Electrochemistry• Electrochemistry is the study of electron
charge transfer processes at an electrode-solution interface.
Ox + ne� � Red
A–Be–
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Electron TransferFe3+
Fe2+
H
H+ H+
He–
e–
Solution
Fe3+
Fe2+
e–
Oxidation
Solution
Fe3+
Fe2+
e–
Electrode
Reduction
Electrode
Fe3+ + e� ! Fe2+ Fe2+ � e� ! Fe3+
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More Examples
Solution
2Cl–2e–
Electrode
Cl2
Producing Chlorine Gas
2Cl� � 2e� ! Cl2
Fe� 2e� ! Fe2+
CorrosionSolution
Fe2+2e–
Iron (Fe)
Solution
Cu2+2e–
ElectrodeCu layer
Cu Deposit growth
Copper Electroplating
Cu2+ + 2e� ! Cu
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Electrochemical (Galvanic) Cell
e–
e–
e–A
e–
Electron Flow
Cathode AnodeHigh Potential
Low Potential
Reduction reaction induces positive potential on electrode relative to solution
Oxidation reaction induces negative potential on electrode relative to solution
A
A
AA
B
B
B B
B
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Electron Transfer at Electrodes Electrode
EF
Solution
(0 eV)
e–
A + e– → A–REDUCTION
Electrode
EF
Solution(0 eV)
e–
B – e– → B+OXIDATIONMetal Electrode
Fermi Level EF
Chemical Species in Solution
Pot
entia
l (eV
)
Vacuum Level (0 eV)
Lowest vacant MO
Occupied MO
Empty States
Filled States
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Electron Transfer at Electrodes
Metal
EFPot
entia
l (eV
)
Vacuum Level (0 eV)
Electron Work
Function
Metal Work Function (eV)
Silver 4.26
Mercury 4.49
Copper 4.65
Gold 5.1
Platinum 5.65
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Electron Transfer at Electrodes
Metal
EF (eV)
Pote
ntia
l (eV
)
Vacuum Level (0 eV)
Redox Species in Solution
Lowest vacant MO
Occupied MO
Pt
-4.0
-4.5
-5.0
Au
Cu
Ag
-5.5
The work function (hence EF value) varies from metal to metal
Silver and Copper Electrodes more likely to Reduce the Species
than Gold or Platinum.
A Platinum Electrode is more likely to Oxidise the Species than
Gold, Copper or Silver.
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Electrode ReactionsNegative Charge
on Electrode
+
-
-
--
-+
+
++
++
+
---
M(s)
Metal electrode M(s) dipped into solution containing corresponding metal ions Mz+(soln)
-
-
--
-+
+
+
+
+
M(s)M(s) → Mz+(soln) + ze–
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Electrode Potential
M+
–
–
–
–
–
+
+
+
+
M+
M+
M+
–
–
–
M+
M+
M+
M+
M+
M+
M+
M+
M+
–
–
–
–
–
–
–
M1(s) ! M1+(sol.) + e� M2
+(sol.) ! M2(s)� e�
[(EM1 � �s)� (EM2 � �s)] = (EM1 � EM2)
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Daniell Cell• Electrode reactions:
• The salt bridge prevents Cu2+ ions going directly to the Zn electrode to pick up free electrons. ‣ This would short-circuit
the battery. ‣ (A porous ceramic usually
replaces the salt bridge)
E = 1.12 V
NaCl Saline Bridge
Copper (cathode)
Zinc (anode)
+ –
CuSO4 soln. ZnSO4 soln.
2e–
Zn2+Cu2+
Cu2+ + 2e� ! Cu
Zn ! Zn2+ + 2e�
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Standard Hydrogen Electrode (SHE)
H2 (1 atm.)
Pt Electrode 2H+
2e- →
For the Standard State ([H+] = 1M, H2 gas at 1 atm, T = 298K) we define: EoH2 / H+, ox. = EoH+ / H2 , red. ≡ 0 V
H2 (1 bar) – 2e– → 2H+ (aH+ = 1)
SHE Half-Cell: Oxidation Reaction at Platinum Electrode (Anode)
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E = Eo + 2.303 RTnF
log10 ( aOx
aR )
The Nernst Equation• Describes how the cell E.M.F. E depends on the standard
potential of a redox couple and on the concentrations of the oxidising and reducing species:
• Given the half-cell reaction: Ox + ne– ⟶ R
the Nernst equation gives:
• Activities aOx and aR are equal to concentrations [Ox], [R] for dilute solutions.
E = Eo + RTnF
ln ( aOx
aR )
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• At the standard temperature (T=25℃) and a single electron reaction the equation simplifies to:
• If reduced species R is a metal electrode it has a constant conc. (aR = 1) and so:
• Similarly if the electrode is the oxidised species:
The Nernst Equation
E = Eo + 0.059 log10 ( [Ox][R] )
E = Eo + 0.059 log10[Ox]
E = Eo −0.059 log10[R]
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Daniell Cell Revisited
Copper (cathode)
Zinc (anode)
+ -
CuSO4 soln.
ZnSO4 soln.Zn2+Cu2+
i 2e-
Cu2+ + 2e� ! Cu
Spontaneous Reaction
Ox + ne� ! R
ECu = 0.3419 + 0.059 log10[Cu2+]
Zn ! Zn2+ + 2e�
Spontaneous Reaction
R ! Ox + ne�
EZn = 0.7618 + 0.059 log10[Zn2+]
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Free Electrons in a MetalFermi Level Energy Levels
occupied by Electrons
Unoccupied Energy Levels
+– Positive Potential
Negative Potential
e–e–e–e–
Current
The Potential Energy of the Electron Energy Levels can be increased or lowered by applying a Negative or Positive Potential.
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Voltage Control of Redox ReactionsApply –ve Potential to Electrode
Apply +ve Potential to Electrode
Electrode
Fermi Level EF
Solution
Pote
ntial
(eV)
Vacuum Level (0 eV)
Lowest vacant MO
Occupied MO
e–
A + e– → A–REDUCTION
ElectrodeEF
SolutionVacuum Level (0 eV)
Electrode
Fermi Level EF
Solution
Pote
ntial
(eV)
Vacuum Level (0 eV)
e–
A – e– → A+OXIDATION
ElectrodeEF
SolutionVacuum Level (0 eV)
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Potential-Current Curve: Butler-Volmer Equation
I
(E–Eo)
IOx
IR
+ve
–ve
Anodic
Cathodic
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Cyclic Voltammetry
(E–Eo) Volt
Curre
nt
Ox + e– ⟶ R
R – e– ⟶ Ox
(+I)
0.0 -0.1 -0.2+0.1+0.2
Cathodic Current
Anodic Current
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Electrode (Surface) Interactions
• Mass Transfer involves: Diffusion of Ox and R down Concentration Gradients.
ne– Electron Transfer
R(surface)
Mass Transfer
Adsorption
Desorpt
ion
DesorptionAdsorption R(bulk)
Ox(bulk)Ox(surface)
Mass Transfer Diffusion Layer
Thickness δ
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Standard Reduction Potentials with a Platinum Electrode
Platinum ~+1V
+0.77 V
Approximate Potential for Zero Current (vs. SHE)
Fe3+ + e → Fe2+
2H+ + 2e → H2
Sn4+ + 2e → Sn2+
Ni2+ + 2e → Ni
+0.00 V
+0.15 V
–0.25 V
SHE
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Amperometric Currents at a Platinum Electrode
~ +1+0.77 +0.15 -0.25
0Potential (Volts vs. SHE)
Cur
rent
Fe3+ + e → Fe2+
Sn4+ + 2e → Sn2+
Ni2+ + 2e → Ni
Reduction Peaks
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Standard Reduction Potentials with a Gold Electrode
Cu2+ + 2e ↔ Cu
Gold ~+0.1V
+0.77 V
Approximate Potential for Zero Current (vs. SHE)
Fe3+ + e ↔ Fe2+
0
+0.34 V
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Amperometric Currents at a Gold Electrode
+0.77+0.340
Potential (Volts vs. SHE)
Cur
rent
Fe2+ - e → Fe3+
Cu - 2e → Cu2+
Oxidation Peaks
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Three-Electrode Electrochemical Cell
−
+
I
WE
CE
RE
WE: Working (indicating, sensing) electrode RE: Reference Electrode CE: Counter (auxiliary) electrode
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Three-Electrode Cell• WE: Ideally polarized electrode ‣ No Faradaic reaction current over the working range
of potentials (Pt, Au?)
• RE: Non-polarisable electrode ‣ Current flow is zero or small currents do not cause a
potential difference (Ag/AgCl)
• CE: Should not affect the reaction at WE ‣ Non-polarisable and very large so current does not
cause a potential difference or limit current
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Reference electrodesPlatinum Wire
acts as Indicator
Electrode that responds to
[Fe2+]/[Fe3+]
Cathode: Fe3+ + e- ↔ Fe2+
Silver Chloride
Anode: Ag + Cl- ↔ AgCl + e-
Salt Bridge
Fe2+ , Fe3+
+–
Saturated KCl solution
Solid KCl
Silver Wire
Ecell =
⇢0.771� 0.059 log10
✓[Fe2+]
[Fe3+]
◆���0.222� 0.059 log10[Cl
�]
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Silver Chloride
Salt Bridge
Fe2+ , Fe3+
+–
Saturated KCl
solution
Solid KCl
Silver Wire
Platinum Wire
Reference Electrode
Reference Electrode: [Cl–] is constant (saturated)
Potential of the Cell only depends on [Fe2+] & [Fe3+]
Ecell =
⇢0.771� 0.059 log10
✓[Fe2+]
[Fe3+]
◆���0.222� 0.059 log10[Cl
�]
Reference electrodes 29
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Ag-AgCl Reference Electrode
AgCl(s) + e– ↔ Ag(s) + Cl–
Eo = 0.22233 V
Air Inlet
Ag Wire (bent into a
Loop)
AgCl Paste
Aqueous solution saturated with KCl
and AgCl
Solid KCl plus some AgClPorous Plug for contact
with External Solution (salt bridge)
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Liquid Junction Potential• Occurs whenever dissimilar electrolyte solutions are in
contact. ‣ Develops at solution interface (Salt Bridge) ‣ Small potential (a few millivolts) ‣ Fundamental limitation on the accuracy
of potentiometric measurements.Different ion mobility results
in charge separation
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Ion Selective Electrodes• ISE respond selectively to one ion
• Contains a thin membrane capable of allowing only the desired ion to bind or to permeate through it
• Sensing does not involve a redox process.
• Electrode Potential defined by Nernst Equation:
• Where [A+] is the activity (conc.) of the ion analyte and n is the charge of the analyte
E = Eo +0.059
nlog10[A
+]
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pH Electrode
Glass sensing membrane
Internal solution: HCl (pH = 7) with KCl/AgCl (saturated)
Internal sensing
electrode: Ag/AgCl
Reference electrode: Ag/AgCl
Reference solution: KCl/AgCl (saturated)
Output voltagedifference between
sensing and reference electrodes
Liquid junction (frit) to measured solution
• Potential generated by H+ difference across glass membrane
• High resistance sensor - needs very high input impedance for instrumentation.
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pH Electrode - Glass Membrane• The outer and inner glass surfaces ‘swell’ to
form a gel as they absorb water.
• The surfaces are in contact with [H+].
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pH Electrode - Glass Membrane• H+ diffuse into glass membrane and replace Na+ in hydrated
gel region.
• There is an ion-exchange equilibrium between H+ and Na+
• Selective to H+ - only ion to bind significantly to the glass gel.
E = constant� �(0.059)pH
Charge is slowly carried by migration of Na+ across glass membrane
(high resistance)
Potential is determined by the [H+] in the external solution.
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pH Electrode OutputE = constant� �(0.059)pH
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