Chapter 22 An Introduction to Electroanalytical Chemistry

1)

* . .

* Electrochemistry (Ionics and Electrodics)

* Babylonians as early as 500 BC used the galvanic cell 50

Knig Bagdad

The formations of electrochemistry are to be found in the late 18th -century investigation by Galvani and Volta.

Galvani frog leg , 1786 ( )

Alecssandro Volta

Volta Cell (battery) led to the beginning of physical electrochemistry.

Ostwald and his associates developed the electrochemical techniques and theories quite rapidly.

William Nicholson (1753-1815) and Anthony Carlisle(1768-1840) used a pile to electrolyze water and various solutions.

Jones Berzelius (1779-1848) and William Hisinger (1766-1852) at 1803 electrolyze salts.

Davy(1778-1848) and Faraday(1791-1867) invented electric motor, generator, and the transformer. Fused salt electrolysis (1807).

In 1805, Christia Grotthus (1785-1822) explains why electrolysis products appear only at the electrodes.

Determination of transference numbers were beginning to open up.

Henry Cavendish (1731-1810) explains relation conductances of water and salt solutions as early as 1776.

In 1847 Eben Horsfoed (1818-1893) studied the effects of electrode polarization.

Friedrich Kohlrausch (1840-1910) in 1869 introduce techniques for the ac method.

Savante Arrehenius (1859-1927) developed his theory of electrolytic dissociation.

In the period 1853 to 1859 Johann Hittorf (1824-1914); the analysis, after electrolysis, of the solutions around one or both of the electrodes.

-o values has been subjected to considerable theoretical study and correction, notable by Peter Debye (1884-1966) and Erich Huckel (1896-1980) and by Lars Onsagar(1903-1976)

Electrical energy - Basic were the investigations of these relationships carried out by James Joule(1818-1889)

Ostwald, Father of physical Chemistry(1853-1932), 1884, Professor of Chemistry

He interested in the emfs of the cells, (He thought that it should be possible to employ a DME as the basis of a system of absolute potentials) factors that govern the emf of cells

Walther Nernst (1864-1941) : assistants of Ostwald in 1889. Nernst's announcement of the relationship that is the basis of electorlytic potentiometry - Nernst equation.

Theodore Richards studied with Ostwald in 1894.

Major contributions to fields such as precise coulometry.

Max Le Blanc(1865-1943) - Studied decomposition voltages of solutions at Pt electrodes. In HCl, halogen can replace oxygen as an anodic product. Separately determine the anodic and the cathodic potentials.

Wilhelm Bottger (1871-1949) used the hydrogen electrode to study the potentiometric titration curves of acids and bases.

Standardization of Potentials. The need for standard electrode potentials increase interest in potentiometry.

In 1890 Ostwald introduced the calomel electrode.

Nernst chose the NHE, assigning to it a potential of zero.

Gilbert Lewis(1875-1946) - The concept of ionic activity in 1907.

Lewis used this concept to set up a listing of electrode potentials based on the standard hydrogen electrode. In 1953, international agreement confirmed the hydrogen scale.

-Electrolytical producing Aluminum is quite expensive by chemical means.

-Cryolite - Alumina electrolytic process in 1886 by C.M.Hall in the U.S.A and by P. Herault in France.

Humphry Davy - potassium, sodium, and other reactive metals in 1807. Fused-salt electrochemistry - Aluminum production is representative

Synthesis of inorganic compounds by electrochemical means contains more than 4,000

Bromine winning form brine.

Faraday, Schoenbein & Kolbe as the founders

ex) - Electroreduction of nitrobenzene - dye stuffs precursors.

- Controlled - potentital electrolysis - Fritz Haber introduce it.

- Various alkaloids are formed in living system.

- Large-scale, adiponitrile - nylon intermediate.

- Modern electrodeless sensing system.

- Copper - electrorefining

- Electrochemical Machining involves the controlled removal of metal

A basics requirement is the ability to measure electrical quantities such as potential, current, & resistance or conductance.

Oersted's 1820 discovery of the deflection of a magnetic needle by an adjacent current carrying wire.

William Thomson designed the galvanometer. - Ink-jet printing.

Poggendorf in 1841 - null-point method - Wheatstones.

- Galvanometric pen recorders - 1890.

- X-Y recorder - in 1951.

- Photographic recorder - 1925.

Development such as that of the vacuum triode, cathode ray tube and, transistor and solid-state IC.

Automation

Emeritus professor I. M. Klothoff

He is the scientific ancestor of generations of electroanalytical chemist.

1) Electrogravimetry;

Szabadvay devotes an entire chapter

1st Electrogravimetric Determination - Wolcott

Gibbs (1864) in 1865, C. Luckow(German)

2) Potentiometry;

Ion selective electrodes.

Glass electrode

The 1st potentiometric titration - R. Behrend in 1893.

The 1st monograph, by E. Mueller (1923)

3) Conductometry;

The 1st analytical determination by this method occurred in 1895

(I. M. Kolthoff)

4) Coulometry;

Faraday's law - 1834.

Silver coulometer, other chemical coulometer.

The first major development occurred in 1938, Hungarian, L. Szebeooedy & Z. Somoggi described the titration of HCl with OH - a good example is the technique of spectroelectrochemistry

5) Voltammetry;

The discovery of polarography - DME

- In 1952. stripping voltammetry

- Pulse technique(1958)

- Cyclic voltammetry - diagnosis of electrode reaction(1964)

6) Amperometric titration;

Titration at one indication electrode was first reproved by Heyrovsky & Berezicky in 1929

- I. M. Kolthoff continued intensively.

1) Basic concepts: A Redox rxn. involves transfer of electrons from one species to another.

-An oxidizing agent

-An reducing agent

ex) Fe3+ + Cu+ = Fe2+ + Cu2+ (1)

Oxidizing Reducing

Agent Agent

Fe3+ + e = Fe2+

Cu+ - e = Cu2+

2) Relation Between Chemistry and Electricity

Electrical Measurement

i) Charge is measured in Coulombs, C

- C of single electron: 1.602189210-19C

One mole of electrons has a charge of 9.648456104C

Faraday constant, F

q = nF

Coulombs = (Mole) (Coulombs/Moles)

ii) Current : The quantity of charge flowing each second through a circuit.

1 A = 1 C/1 sec

iii) Voltage, Joule and Free Energy

One J of energy is gained or lost when one coulomb of charge is moved through a potential difference of one volt.

Work = Eg

Joules = VoltsCoulomb

Work = -G at constant T, P reversible chemical reaction

G = - Work = - Eq

G = - nFE

Ohm's law

I = E/R

Power (P)

P = Work/s = (Eq)/s = E (q/s) = EI [W]

4) Galvanic Cells

i) Half reactions

ii) Anode : The electrode at which oxidation occurs.

iii) Cathode : The electrode at which Reduction occurs.

iv) Salt bridge

We adopt the convention that the left-hand electrode of each cell is connected to the negative input terminal of the meter.

left-hand side : oxidation electrode

v) Line notation

5) Standard potentials

i) Standard reduction potential.

ii) SHE NHE Potential assign zero

- Using the Nernst equation.

Anode : H2(g) ---> 2H+ + 2e-

Cathode : Cd2+ + 2e- ---> Cd(s) Eo = -0.402

-----------------------------

Net : Cd2+ + H2(g) ---> Cd(s) + 2H+

E(cell) = Eo(cell) - (0.05916/2) log ([H+]/[Cd2+]PH2)

pH --> 1 if E=0 --> [H+] = 1.610-7

Solubility Product: [Ag+] = (Ksp([AgCl])/[Cl-] = [(1.810-10)/0.0334]

= 5.410-9M

(K --> Eo)

E(cell) = Eo-(0.0591/n) log Q (at any time)

O = Eo-(0.0591/n) log K (at eq)

(0.0591/n) log K = Eo or K = 10nEo/0.05916 (at 25 C)

ex) FeCO3(s) + 2e- --> Fe(s) + CO32- Eo = - 0.0756 V

Fe(s) --> Fe2+ + 2e- Eo = - 0.440 V

-----------------------------------------------------------------

FeCO3(s) --> Fe2+ + CO32-(K=Ksp) Eo = - 0.316V

Ksp = 10(2)(-0.316)/(0.0591) = 210-11

- Using Cells as Chemical probes

a) Eq. between the two half-cells

b) Eq. within each half-cell

1. : Leclanche or dry cell (1.5V)

(ZnCl2. NH4Cl, MnO2, Zn, C)

Anode : Zn(s) ---> Zn2+ + 2e-

Cathode : 2NH4+ + 2e- ---> 2NH3(g) + H2(g) : Carbon

-

2MnO2(s) + H2(g) ---> Mn2O3(s) + H2O(l)

Zn2+(aq) + 2NH3(g) + 2Cl- ---> Zn(NH3)2Cl2(s)

2MnO2(s) + 2NH4Cl + Zn(s) ---> Zn(NH3)2Cl2 + H2O + Mn2O3(s)

2. . (1.54V)

Zn(s) + 2OH- ---> ZnO(s) + H2O(l) + 2e-

2MnO2(s) + H2O(l) + 2e- ---> Mn2O3(s) + 2OH- -->

3. (1.35V)

Anode : Zn(s) + 2OH-(aq) ---> ZnO(s) + H2O(l) + 2e-

Cathode : HgO(s) + H2O(l) + 2e- ---> Hg(l) +2OH-(aq)

4.

- : + (10) in C-H2SO4

Anode : Pb(s) + SO42-(aq) ---> PbSO4(s) + 2e-

Cathode : NiOOH(s) + H2O(l) + e- ---> Ni(OH)2(s) + OH-(aq)

5. Fuel Cell

Anode (H2): 2H2(g) + 4OH-(aq) ---> 4H2O(g)

Cathode (O2 ): O2(g) + 2H2O(l) + 4e- ---> 4OH-(aq)

70 - 140 C

Output ---> 0.9V

Fused salt

Anode : 2Cl- ---> Cl2(g) + 2e- Eo = - 1.36V

Cathode : Na+ + e- ---> Na(s) Eo = - 2.71V

---------------------------------

Net 2Cl- + 2Na+ ---> 2Na(s) + Cl2(g)

1) Aluminum

: 3Na(s) + AlCl3(s) ---> Al(s) + 3NaCl(s)

: Anode: O2 Cathode: Al

2) Na : Humphrey Davy

NaOH ---> Na

3) Cl2 and NaOH

NaCl ---> Cl2 + 2e-

Na+ + e- + Hg ----> Na(Hg)

2Na(Hg) + 2H2O(l) ---> 2NaOH(aq) + H2(g) + Hg(l)

{ Corrosion }

Fe ----> Fe2O3.H2O 25%

2Fe(s) + 3O2(g) ---> 2Fe2O3

4Cu(s) + O2(g) ---> 2Cu2O(s)

4Al(s) + 3O2(g) ---> 2Al2O3(s)

Anode : M(s) ---> Mn+ + ne-

Cathode : 2H+(aq) +2e- ---> H2(g)

2H2O(g) + 2e- ---> 2OH-(aq) +H2(g)

O2(g) + 2H2O(l) + 4e- ---> 4OH-(aq)

()

Anode : Fe(s) ---> Fe2+(aq) + 2e-

Cathode : 2H2O(l) +2e- ---> 2OH-(aq)+ H2(g)

--------------------------------------------------

Net : Fe(s) + 2H2O(l) ---> Fe2+(aq) + 2OH-(aq) + H2(g)

Fe(OH)2(g) - Fe2O3

+ ---> 100 .

Anode : 2[Fe(s) ---> Fe2+(aq) + 2e-]

Cathode : O2(g) + 2H2O(l) + 4e- ---> 4OH-(aq)

---------------------------------------------------

Net : 2Fe(s) + O2(g) + 2H2O(l) ---> 2Fe(OH)2(s)

6Fe(OH)2(s) + O2(s) ---> 2H2O(l) + 2Fe2O3H2O

---> 2Fe3O4H2O(s) + 4H2O(l) .

Fe3O4H2O(s) ---> H2O(l) + Fe3O4(s) .

Cl-

- Anodic inhibition :

ex) 2Fe(s) + 2Na2CrO4(aq) + 2H2O ---> Fe2O3 + Cr2O3 + 4NaOH(aq)

Na2CrO4 .

- Cathodic protection :

. ex)

1.1 Electrochemical method

Analytical balance

Hot plates

Hume hoods

Ovens

pH meter

UV/VIS & IR - spectrophotometer : 50 %

AA - spectrophotometer : 30 %

Polarographic analyzer : 12 %

Ion - selective electrode : 30 %

: .

A)

Inexpensive

Specific for a particular chemical form

Concentration activity

B) Electrochemical method classification

Electrochemical method :

C) Experimental system

Electrolyte

Detector (electrode)

Circuit

D) Electrochemical method

Potentiometric All others (voltammetry, coulometry, conductometry, etc.)

Potentiometry : J. Willard Gibbs Nernst( )

System

All others : voltage or current voltage system monitor

22A Electrochemical cells

- 2

(1) Electrode

- Anode oxidation

- Cathode reduction

1) Working and indicator electrode : A reaction take place.

2) Reference electrode : constant

potential

3) Counter electrode : internal polarization

E(working electrode) = (Ecell - iRcell - Epolarization)

- Electrode

- Electrolyte

(1)

* Galvanic cell : chemical En. electrical En.

* Electrolytic cell : electrical En. chemical En.

Galvanic cell : Zn Zn2+ + 2e-

Cu2+ + 2e- Cu

(2)

Zn / Zn2+(aZn2+)Cu2+(a Cu2+) / Cu

Left hand electrode : negative pole of cell

oxidation process occurs

Zn Zn2+ + 2e

Right hand electrode : Reduction process

Cu2+ + 2e Cu

*

Pt, H2(p=1atom) / H+(0.1M), Cl-(0.1M), AgCl(satd) / Ag

22A-5 Solution Structure

Fig. 22-3 Electrical double layer

1) A compact inner layer (d0 to d1): the potential decreases linearly with distance from the electrode surface.

2) A diffuse layer(d1 to d2): the potential decrease is exponential.

22A-6 Faradaic and Nonfaradaic Currents

Faradaic process : - Faraday's law ( ) .

2) Non-faradaic process : - condenser

3) Charging current : non-faradaic process

EA EB . charging current.

22B Potentials in Electrochemical Cells

Ecell = Eright - Eleft + Elj

22B1. Thermodynamic cell potential

Nernst Equation thermodynamic relationship potentiometric measurement

G = H - TS = E + PV -TS

From Vant Hoff reaction equation = dE - TdS -SdT + PdV + VdP

Ion , error

A) Standard electrode potential: , l .

B) Formal electrode potential: (, , , ) swift formal .

Formal potential

(, )

@ Effect of complexation on the electrode potential.

->

ex) Zn / Zn2+ // Cu2+ / Cu CuSO4 solution EDTA

-> Cu2+ + EDTA4- Cu EDTA2-

22B-2 Liquid - Junction Potentials

* ionic solution

A) Diffusion potential

Liquid junction potentials, different mobilization & concentrations of ions in electrolytes in contact.

B) Donnan potential

1 ions .

salt bridge, porous glass Kl : salt bridge.

C) Liquid junction

Transport number

Charge

Activity of the ions forming the junctions

1) Salt bridge

2) Cracked glass bead

3) Ceramic frit

4) Sleeve

5) Gauntly or asbestos fiber, wick

6) Platinum wire

7) Cellulose pulp

8) Glass frit

9) Cellophane

10) Fine capillary drip

:

Liquid junction .

22E Currents in an electrochemical cell

Cell potential ( )

Thermodynamic cell potential

Liquid junction potential

Ohmic potential

Polarization potential { , }

Thermodynamic cell potential

Liquid junction potential

Polarization potential concentration polarization

overvoltage(kinetic polarization)

IR .

Fig. 22-6 Curves for an ideal(a) polarized (b) nonpolarized electrodes

: , , ,

, , , , .

1) Ideally non-polarized electrode.

- (condenser) . ,

2) Ideally polarized electrode.

- .

, KCl (polarography )

K+ + e K (amalgam), 2Hg+ Hg22+ + 2e( )

2Cl- Cl2 + 2e ( ), 2H2O + 2e- H2 + 2OH- ( .)

* Capacitance of an electrode

condenser

3) Depolarizer : polarized electrode

4) Reversible : .

:

* : Concentration polarization, charge transfer polarization, kinetic polarization,( )

* : bulk

* : overvoltage

Fig. 22-8

Electrode Surface layer Solution bulk

*

Change transfer limiting step : activation polarization

Mass transfer limiting step : concentration polarization

Chemical reaction or adsorption

5) Overvoltage(overpotential)

a) Activation overvoltage;

Slow electron transfer high activation En

b) Resistance overvoltage;

iR drop , .

: adherent layer

c) Concentration overvoltage;

Electrode vicinity concentration

polarized : stirring

3 .

overvoltage

= - e

: Potential difference across the interface

e : Potential difference across the interface at equilibrium

* Non-polarizable electrode :

Reference electrode non polarizable electrode .

1) Activation overvoltage : slow electron transfer

2) Resistance overvoltage iR drop

3) Concentration overvoltage

= A + C + R

Modes of Electrochemical Mass Transport

Migration, convection, diffusion 3 .

1) Migration : electrical gradient ( )

: or indifferent electrolytes .

2) Convection : gross physical movement

3) Diffusion : most widely studied