Introduction to Electrochemistry · 2019-06-30 · U-Tokyo Special Lectures Biosensors and...

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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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

Electrochemistry

• Electrochemistry is the study of electron charge transfer processes at an electrode-solution interface.

3

Ox + ne� � Red

A–B

e–

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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+

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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.

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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–

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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)

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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�

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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+

]

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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.

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

Potential-Current Curve: Butler-Volmer Equation

19

I

(E-Eo)

IOx

IR

+ve

–ve

Anodic

Cathodic

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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)

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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.

Stewart SmithBiosensors and InstrumentationU-Tokyo Special Lectures

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

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