Introduction to Electrochemistry - University of Edinburgh · Introduction to Electrochemistry Lecture 4 1. Biosensors and Instrumentation Stewart Smith Beijing University of Posts

Stewart SmithBiosensors and InstrumentationBeijing University of Posts and

Telecommunications 2019

Introduction to Electrochemistry

Lecture 4

1



Summary

• Redox reactions

• Standard electrode potential

• Control of electrode reactions

• 3-Electrode cell and Reference electrodes

• Ion sensitive electrodes

2



Electrochemistry• Electrochemistry is the study of electron

charge transfer processes at an electrode-solution interface.

Ox + ne� � Red

A–Be–

3



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+

4



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

5



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

6



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

7



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

8



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.

9



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–

10



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)

11



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�

12



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)

13



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 )

14



• 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]

15



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+]

16



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.

17



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)

18



Potential-Current Curve: Butler-Volmer Equation

I

(E–Eo)

IOx

IR

+ve

–ve

Anodic

Cathodic

19



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

20



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

21



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

22



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

23



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

24



Amperometric Currents at a Gold Electrode

+0.77+0.340

Potential (Volts vs. SHE)

Cur

rent

Fe2+ - e → Fe3+

Cu - 2e → Cu2+

Oxidation Peaks

25



Three-Electrode Electrochemical Cell

−

+

I

WE

CE

RE

WE: Working (indicating, sensing) electrode RE: Reference Electrode CE: Counter (auxiliary) electrode

26



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



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

�]

28



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



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)

30



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

31



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

+]

32



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.

33



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



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.

35



pH Electrode OutputE = constant� �(0.059)pH

36

Documents

Introduction to Electrochemistry - University of Edinburgh · Introduction to Electrochemistry Lecture 4 1. Biosensors and Instrumentation Stewart Smith Beijing University of Posts