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Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
Introduction to Electrochemistry
Lecture 2
1
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
Summary
• Redox reactions
• Standard electrode potential
• Control of electrode reactions
• 3-Electrode cell and Reference electrodes
• Ion sensitive electrodes
2
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Electrochemistry
• Electrochemistry is the study of electron charge transfer processes at an electrode-solution interface.
3
Ox + ne� � Red
A–B
e–
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Electron Transfer4
Fe3+
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 Examples5
Solution
2Cl–2e–
Electrode
Cl2
Producing Chlorine Gas
2Cl� � 2e� ! Cl2
Fe� 2e� ! Fe2+
CorrosionSolution
Fe2+2e–
Iron (Fe)
Solution
Cu2+2e–
Electrode
Cu layer
Cu Deposit growth
Copper Electroplating
Cu2+ + 2e� ! Cu
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Electrochemical (Galvanic) Cell6
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 7
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 8
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 9
MetalEF (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 Reactions10
Negative 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 Potential11
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)
12
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)13
H2 (1 atm.)
Pt Electrode2H+
2e- →
For the Standard State ([H+] = 1M, H2 gas at 1 atm, T = 298K)
we define: EoH2 / H+, ox. = Eo
H+ / H2 , red. ≡ 0 V
H2 (1 bar) – 2e– → 2H+ (aH+ = 1)
SHE Half-Cell:
Oxidation Reaction at Platinum Electrode (Anode)
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
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.
14
E = Eo +RT
nFln
✓aOx
aR
◆
E = Eo
+ 2.303RT
nFlog10
✓aOx
aR
◆
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
The Nernst Equation
• 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:
15
E = Eo + 0.059 log10
✓[Ox]
[R]
◆
E = Eo + 0.059 log10[Ox]
E = Eo � 0.059 log10[R]
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
Daniell Cell Revisited16
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 Metal17
Fermi 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 Reactions18
Metal Electrode
Fermi Level EF
Solution
Pot
entia
l (eV
)
(0 eV)Vacuum Level
Lowest vacant MO
Occupied MO
Electrode
EF
Solution
(0 eV)Vacuum Level
e
A + e– → A–
Apply –ve Potential to Electrode
REDUCTION
Metal Electrode
EF
Solution
Pot
entia
l (eV
)
(0 eV)
Electrode
EF
Solution
(0 eV)
eA – e– → A+
Apply +ve Potential to Electrode
OXIDATION
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Potential-Current Curve: Butler-Volmer Equation
19
I
(E-Eo)
IOx
IR
+ve
–ve
Anodic
Cathodic
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Cyclic Voltammetry20
(E-Eo) Volts
Cur
rent
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) Interactions21
ne– Electron Transfer
R(surface)
Mass Transfer
Adsorption
Desorpt
ion
DesorptionAdsorption R(bulk)
Ox(bulk)Ox(surface)
Mass Transfer
• Mass Transfer involves: Diffusion of Ox and R down Concentration Gradients.
Diffusion Layer Thickness δ
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Standard Reduction Potentials with a Platinum Electrode
22
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
23
~ +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
24
Cu2+ + 2e ↔ Cu
Gold ~+0.1V
+0.77 V
Approximate Potential for Zero Current (vs. SHE)
Fe3+ + e ↔ Fe2+
0
+0.34 V
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
Amperometric Currents at a Gold Electrode
25
+0.77+0.340Potential (Volts vs. SHE)
Cur
rent
Fe2+ - e → Fe3+
Cu - 2e → Cu2+
Oxidation Peaks
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Three-Electrode Electrochemical Cell26
−
+
I
WE
CE
RE
WE: Working (indicating, sensing) electrode
RE: Reference Electrode
CE: Counter (auxiliary) electrode
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
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
27
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Reference electrodes28
Platinum 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
✓[Fe
2+]
[Fe
3+]
◆���0.222� 0.059 log10[Cl
�]
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
29
Silver Chloride
Salt Bridge
Fe2+ , Fe3+
+-
Saturated KCl solution
Solid KCl
Silver Wire Platinum
Wire
Reference Electrode
ReferenceElectrode:[Cl-]isconstant(saturated)
Potential of the Cell only depends on [Fe2+] & [Fe3+]
Ecell =
⇢0.771� 0.059 log10
✓[Fe
2+]
[Fe
3+]
◆���0.222� 0.059 log10[Cl
�]
Reference electrodes
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
Ag-AgCl Reference Electrode30
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.
31
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
32
E = Eo
+
0.059
nlog10[A
+]
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
pH Electrode33
Glass sensingmembrane
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 referenceelectrodes
Liquid junction (frit) tomeasured 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+].
34
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
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.
35
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.
Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures
pH Electrode Output36
E = constant� �(0.059)pH