Chapter 7 Electrochemistry

§7.8 Electrode potential

Daniell cell

How does electrode potential establish?

Zinc metal

Copper metal

ZnSO4 solution

CuSO4 solution

Zn2+

Zn2+

Zn2+

Zn2+

Cu2+

Cu2+

Cu2+

Cu2+

Porous partition diagram

electromotive forces

potential difference

1) Metal-metal interface: contact potential + + + + + +

Cu

Zn

7.8.1. Interfacial charge and electrode potential

KCl solution HCl solution

2) Liquid-liquid interface: Liquid junction is the interface between two miscible electrolyte solutions.

liquid junction potential, liquid potential, diffusion potential

3) Liquid-metal:

Cu2+ + 2e Cu

+

+

+

+

+

+

+

+

+

+

++

+

+

+

+ +

+

exchange current, electrode potential

7.8.2. Models of electric double layer

1) Holmholtz double layer (1853)

Compact double layer

++++++++

0

d

E

2) Gouy-Chappman layer (1910, 1913)

Diffuse double layer

++++++++

0d

E

3) Stern double layer (1924)

++++++++

0

d

E

7.8.3 Electromotive forces and relative electrode potential

Cu(s)Zn(s)ZnSO4(m1)CuSO4(m2)Cu (s)

E = c + + j + +

anode cathode

E = c + (l,1- ) + ( +- l,2)

= +- + (c+ l,1- l,2) When emf of a cell was measured, we , in fact, measured

the potential difference between the two electrodes.

E

Can the absolute potential of electrode be unmeasured?

Absolute potential

lm

Only the difference between two electrodes, i.e.,

electromotive of the cell E = + can be measured.

potentiometerarbitrary reference

(2) Normal/Standard Hydrogen Electrode (NHE/SHE)

In 1953, IUPAC defined normal hydrogen electrode (NHE) as the reference for measurement of electrode potential. IUPAC conventions

acidic solution with activity of H+ equals to 1.

platinized platinum foil electrode

pure hydrogen gas at standard pressure

1

H H H1.0mol kg , 1, 1m a

definition

H+/H2 = 0.000000 V.

(3) standard electrode potential

The potential of other electrode can be obtained by

combination of NHE and any other unknown electrode into

an electrochemical cell with NHE serving as negative

electrode and the unknown electrode as positive electrode:

- NHE || unknown electrode +

The sign and the value of the emf of the cell is thus the sign and value of the potential of the unknown electrode.

All standard electrode potentials are reduction potentials.

Cf. Levine, p. 431-435

NHE Cu2+ (a=0.1)Cu

E = 0.342 V

the electrode potential of the Cu|Cu2+ (a=0.1) electrode at

pressure p and temperature T is thus

2Cu / Cu0.342V

Example

Because the unknown electrode is always arranged as positive electrode, the electrode reaction is, therefore, written in reduction form.

(reduction)(standard) electrode potential

Cu2+ + 2e Cu

SHE||Cu2+ (a=0.1)|Cu

Cell reaction: H2(g, p) + Cu2+(a) = 2H+ (a=1) + Cu

22

2CuH

H Cu

lna aRT

E EnF a a

y

(4) Nernst equation for electrode

2+ + 2+2+ +2 Cu /Cu H /H2

Cu /Cu H /H Culn

RTa

nF y y

2+ 2+2+Cu /CuCu /Cu Cu

lnRT

anF

y

ox

red

lnaRT

nF a y red

ox

lnaRT

nF a y

7.8.4 Reference electrode

Problems with NHE (primary standard):

1)The platinized platinum electrode is easily poisoned by adsorption of impurities from the solution and the gas.

2) An elaborate purification is required to purify the hydrogen before it is passed through the cell.

3) Changes in barometric pressure or in the depth of immersion of the electrode in the solution produce a small variation in the potential of the electrode.

4) The preparation and the maintenance of the unit activity solution are both much complicated.

The calomel electrode:

Hg(l)Hg2Cl2(s)KCl (m)

Some electrodes with stable potential usually used as the secondary standard, named as reference electrode

saturated calomel electrode (SCE)

(T)/V ＝ 0.2412 - 6.61 10-4 (T/ -25) - 1.7 ℃

10-6 (T/ -25)℃ 2 - 9 10-10 (T/ -25)℃ 3

HgHg2Cl2

past

mercury-mercurous sulfate electrode:

Hg(l)Hg2SO4(s)SO42(m) ⊖ ＝ +0.640 V

mercury-mercuric oxide electrode: Hg(l)HgO(s)OH(m) ⊖ ＝ +0.098 V

silver-silver chloride electrode: Ag(s)AgCl(s)Cl(m) ⊖ ＝ 0.197 V

⊖(Ag+/Ag) ＝ +0.799 V vs NHE

⊖(Ag+/Ag) ＝ __?___ V vs SCE

Other common reference electrodes

Ag+/Ag0.799 V

0.2412 V

0.000 V

SCE

NHE

7.8.5. liquid junction potential and salt bridge

1) liquid junction potential

The diffusion of ions is irreversible, which destroys the reversibility of the cell.

The value of Ej can reach ca. 30 mV, which is too large

for the measurement of emf.

2) influential factor of Ej

Pt(s), H2(g,p)HCl(m)HCl(m′)H2(g, p), Pt(s)

Ej

On passage of 1 mole of electrons through the cell, t+ mol H+ and t mol Cl pass the boundary

-+

+ -

ClH

H Cl

''ln lnj j

aaG t RT t RT nFE

a a

For uni-univalence electrolytes

'ln)12(

'ln)(

m

m

F

RTt

m

m

F

RTttE j

3) Effects of salt bridge

Every measurement of emf of a cell whose two electrodes require different electrolyte raises the problem of the liquid junction potential.

The problem can be solved either by measuring the junction potential or eliminating it. The salt bridge is often used to connect the two electrode compartments to reduce the junction potential.

4) Salt bridge electrolyte1) does not react with either solution2) transference number of cation and anion is close3) of high concentration.

ions K+ NH4+ Cl NO3

102 /S·m2·mol-1 0.7352 0.734 0.7634 0.7144

t+ of some common salt bridge electrolytes

c/mol·dm-3 0.01 0.05 0.10 0.20

KCl 0.490 0.490 0.490 0.489

NH4Cl 0.491 0.491 0.491 0.491

KNO3 0.508 0.509 0.510 0.512

m

c/mol·dm-3 0 0.2 1.0 2.5 3.5

Ej 28.2 19.95 8.4 3.4 1.1

Concentration-dependence of Ej

Why does salt bridge reduce the junction potential.

+ + + + +- - - - -

5) Effects of salt bridge:

6) Elimination of junction potential