Curs Chimie Cap4

Electrochemistry 89

4. ELECTROCHEMISTRY

Electrochemistry is that part of chemistry which studies the interaction

between matter and electric field. It consists in two parts: ionic electrochemistry

and electrodic chemistry.

Ionic electrochemistry studies the appearance of the ions in solution, the

interactions between them or with other particles, the transport phenomena of ions:

diffusion, migration, conduction.

Electrodic electrochemistry studies what happened at the interface solid –

solution: electrical double layer, electrode potential, galvanic cells, heterogenous

redox reactions, chemical sources of energy, corrosion and protection methods

against corrosion.

Electrochemical processes are important to be studied since electronics

deals with semiconductors, galvanic cells, electrochemical depositions on circuits

etc. This study includes the understanding of the main electrode processes, the

current distribution and reaction phenomena.

4.1. Electrolyte solutions. Electrolytic dissociation

Electrical electrolytes can be divided into:

− 1st degree electric conductors – electronic conductors: metals,

semiconductors, alloys;

− 2nd degree electric conductors – ionic conductors: electrolytes

as solutions or melted electrolytes.

90 GENERAL CHEMISTRY

In the case of electronic conductors the main charge carrier are electrons

and passing an electric current does not lead to chemical modifications in the

material structure.

By definition, electrolytes are chemical species which dissolved in certain

solvents or melted down are able to generate ions. When an electric field is applied

to an electrolyte, the electric current flows, because ions in that particular

environment migrate to the oppositely charged electrodes in contrast to an

electronic conductor where the unique carriers are the electrons.

Electrolytes may be divided into true electrolytes and potential

electrolytes.

A true electrolyte has an ionic structure and forms in solid state ionic

lattice (ions are maintained in fixed positions). By solvating in water, the

electrostatic bonds are broken and free ions appear in solution.

Examples: NaCl, KCl, Al2(SO4)3, CuSO4, NiSO4, ZnSO4

H2O NaCl Na+ + Cl-

H2O

AgNO3 Ag+ + -3NO

A potential electrolyte is a polar covalent molecule. The ions are

generated by reaction with the solvent.

Examples: HCl + H2O H3O+

NH3 + H2O NH4+ + HO-

The process that may occur when a chemical compound is dissolved in a

solvent breaking up into two or more ions forming an ionic conducting solution is

called electrolytic dissociation. This process takes place before the action of

electric field and it is independent on it.

The equilibrium constant which characterizes the electrolytic dissociation

is called dissociation constant Kd.

Types of dissociations:

a) Dissociation of strong acids

H2SO4 2H+ + SO42-

Electrochemistry 91

aSOH

SOHd K

c

ccK =

⋅=

−+

42

24

2

; [177]

The dissociation constant Ka is called the acidity constant

b) Dissociation of bases

Cu(OH)2 Cu2+ + 2HO-

bOHCu

CuHO

d Kc

ccK =

⋅=

+−

2

2

)(

2

; [178]

The dissociation constant Kb is called the basicity constant

c) Salt dissociation

NiSO4 Ni2+ + SO42-

4

24

2

NiSO

SONid c

ccK

−+ ⋅= [179]

Not all the dissolved molecules are always dissociated. The ratio of

dissociated molecules at equilibrium and the total number of molecules in solution

is called degree of dissociation, α.

% ,100molecules dissolvedmolecules ddissociate

×=α [180]

The electrolytes with α < 5 % are weak electrolytes (solutions of 0.1 n or

more diluted). The electrolytes with 5% < α< 50 % are medium electrolytes and

the electrolytes with α > 50 % are strong electrolytes. It is possible to have

electrolytes in which all molecules dissociate (α = 100 %).

Some electrolytes dissociate in several steps. Each step may have different

degree of dissociation:

H3PO4 + H2O H3O+ + H2PO4- big α

H2PO4- + H2O H3O+ + HPO4

2- small α

HPO42- + H2O H3O+ + PO4

3- very small α, the reaction

takes place in special conditions.


d) Dissociation of weak electrolytes

Weak electrolytes could be:

- organic acids CH3COOH, HCN

- inorganic weak acids H2S, H2SO3

- weak bases Al(OH)3, NH4OH

Let us consider the dissociation of a weak acid:

HA H+ + A-

HA

AHd c

ccK

−+ ⋅= [181]

The number of dissociated molecules is equal to the molar

concentration of A- and H+ ions +− = HAcc .

The number of dissolved molecules is equal to the initial molar

concentration of the acid 0HAc

⇒ 00HA

H

HA

A

cc

cc +−

==α [182]

At equilibrium:

)1(000 αα −=−= HAHAHAHA cccc [183]

⇒)1(0

00

ααα−⋅

=HA

HAHAd c

ccK [184]

⇒)1(

02

αα−

= HAd

cK [185]

Equation [185] represents the Ostwald’s Law or Law of dilution

If the dissociation constant Kd for a weak acids is known then it is

possible to calculate the degree of dissociation for some of their concentration:

0

02

24

HA

dHAdd

cKcKK +±−

=α [186]

or knowing Kd and α one can calculate the concentration:

Electrochemistry 93

20 )1(

αα−

= dHA

Kc [187]

e) Water dissociation

Pure water is a weak electrolyte and exhibits a very low electrical

conductivity.

H2O H+ + HO-

OH

HOHd c

ccK

2

−+ ⋅= [188]

OHdHOHcKcc

2⋅=⋅ −+ [189]

The product of hydrogen ions (H+) concentration and hydroxide ions

(HO-) concentrations takes a constant value at a given temperature. It is called

ionic product of water, Kw. 1410−=⋅= −+ HOHw ccK (at 22°C) [190]

In pure water 710−== −+ HOHcc iong/L.

One defines:

pH which is the decimal logarithms of the hydrogen ions

concentration, taken with (-) sign,:

pH = – log +Hc

[191]

The solutions with 0 < pH < 7 are acid solutions ( −+ > HOHcc ), the

solutions with pH = 7 are neutral solutions ( −+ = HOHcc ) and the solutions with

7 < pH < 14 ( −+ < HOHcc ) are basic solutions.

f) Dissociation of strong electrolytes

The strong electrolytes have the dissociation degree close to 1. To this

category belong all the salts, strong acids (HCl, H2SO4), strong bases (NaOH,

KOH).


4.2. Electrical conductivity of the electrolytes

Electrical conductivity represents the substance property to conduct

electric charge (charge mobility). For all conductors the Ohm law’s says: E = RI,

where E represents the applied voltage, R is the ohmic resistance, I is the current

intensity.

1st degree electric conductors

For the 1st degree electric conductors the ohmic resistance is slR ρ

= where

ρ is the conductor resistivity, l is the conductor length and s is the section of the

conductor. The inverse of ohmic resistance is called electrical conductance, G. The

inverse value of the conductor resistivity is called conductivity or specific

conductivity, λ.

GR=

1 [Ω-1] or [S] (Siemens) [192]

λρ=

1 [193]

2nd degree electric conductors

Conductivity, λ, is the conduction capacity of a volume of 1 cm3 of

electrolyte solution when subjected to an electric field equal to 1 V/cm. The

conductivity of the electrolytes is tens or hundred times smaller than the

conductivity of metals. At the same time the electrical conductivity of some weak

electrolytes as water or alcohols is a several orders of magnitude bigger than those

of dielectrics.

It is important to know the electrical conductivity in order to reduce the

electrical energy consumption (using electrolytes with higher electric

conductivity).

Electrochemistry 95

Equivalent conductivity, Λ, represents the conductivity of a volume of

electrolyte which contains ions of one equivalent gram dissociated substance. The

ions are located between the electrodes of a conducting cell while the distance

between the electrodes is 1 cm.

cλ⋅

=Λ1000 [194]

where c is the concentration expressed in ion g/L.

The equivalent conductivity depends on concentration, degree dissociation

and charge mobility.

Kohlrausch's Square Root Law: For strong electrolytes sufficiently

diluted, the molar conductivity Λc of an electrolyte with concentration c is a linear

function of the its square c . If Λ0 represents the molar conductivity at infinite

dilution and A represents a constant which is function on the valence of the salt

then:

cc Α−Λ=Λ 0 [195]

For weak electrolytes:

0ΛΛ

= cα [196]

αα−

=1

2cKd (Ostwald Law)

⇒ )( 00

2

c

cd

cK

Λ−ΛΛΛ

= [197]

Kohlrausch's Law on the independence of migrating ions: The molar

conductivity of an electrolyte equals the sum of the molar conductivities of the

cations and the anions; ν being the number of anions or cations:

−−++ Λ+Λ=Λ νν [198]

It results that in electrolytes the contribution of each ion to the equivalent

conductivity is the same, and does not matter the chemical compound.


4.3. Electrical charged interfaces

Interface represents a surface contact between two phases. Usually it is

thin and defined as a monoatomic layer.

Interphase represents a region with modified properties with respect to

those inside the phase. It is a region of at least two molecular diameters. Interphase

is a region between two phases. In this region the properties are not the same as in

each phase. The properties of this interphase region influenced and even dominate

the electrochemical reactions of the charge transfer.

Figure 24. Schematic representation of the interface and interphase

If the two phases contain charged particles (ions or electrons) than, as a

consequence of the rearrangements of these particles in the interphase region, the

interface is charged: the two phases are charged with equal, but opposite sign,

charges. The interphase region is electrically neutral.

In this way a double electric layer, similarly to a condensator, is formed. This

layer is characterized by a variation of electrical potential, as shown in figure 25. Processes which influences the formation of a double layer:

- agglomeration of the electrons at the metal surface as a consequence

of the fact that in this region the electrons could arrive to the exterior at

a distance larger than atomic ray. Negative charges of the metal attract

the positive ions from the solution or orient the dipoles;

Air

Solution

Interphase Interface

Electrochemistry 97

- specific adsorption of some ionic species of the solution at the

interface followed by the electrostatic attraction of the ions of opposite

sign from the solution;

- oriented dipole adsorption at the interface;

- transfer of the electrical charged particles through the interface

metal/own ions solution system M/Mz+.

a) b)

Figure 25 Schematic representation of the double electric layer (a) and the variation of

electrical potential at the double layer

4.3.1. Internal Potential, External Potential, Surface Potential

Electrode represents a system formed from different phases charged

electrically and being in contact.

Electrode potential is the difference in electrical potential of metal and

solution phases.

A phase conducting electric current is characterized by an internal

potential or Galvani potential, Φ.

The internal potential is the electrostatic potential of one point within the

phase compared to the potential of one point which is placed at infinite in the

δ

Phase I Phase II - - - -

+ + + +

Electrical Potential

Distance

δ


uncharged vacuum. The internal potential is the work necessary to bring the

elementary charge from the infinity to inside the phase.

The internal potential has two distinct components. One is the Volta

potential, Ψ which is determined by the long range coulombic forces near the

electrode, and the surface potential, χ which is determined by the short range

effects of adsorbed ions and oriented water molecules:

Φ = Ψ + χ [199]

The Volta potential can be measured directly while the surface potential

cannot. Thus the Galvani potential can be only measured relative to a reference

electrode.

4.3.2. Equilibrium Electrode Potential

In order to establish the equilibrium conditions in electrochemical systems

M/Mz+ Gugenhein (1929) introduced the electrochemical potential iµ :

iµ = µi + ziFεI [200]

where i represents the component; µi represents the chemical potential, zi

the valence and εi the equilibrium potential of the phase of ith component. F is the

Faraday constant; F = 96500 C/mol.

a) System M/Electrolyte solution

At equilibrium the electrochemical potential of the cations from solution

(s) is equal to the electrochemical potential of the cations from metal (m).

++ = zz MMMS µµ [201]

⇒ Sµi + zFεS = Mµi + zFεM [202]

The equilibrium potential is:

ε = εM – εS [203]

⇒ ε = zF1 (Sµi – Mµi) [204]

But

Electrochemistry 99

µI = µi0 +RT ln ai [205]

where ai is the activity of the ith ions. The activity represents the deviation

of the behaviour of electrolyte from the ideal solution, behaviour characterized by a

total dissociation.

µi0 is the standard chemical potential.

⇒ ε = zF1 (Sµi

0 – Mµi0) + M

i

Si

aa

zFRT ln [206]

For the metal (pure component) Mia =1

⇒ ε = ε0 + Sia

zFRT ln [207]

Equation [207] is called Nernst equation

In this equation: ε is the electrode equilibrium potential

ε0 is the standard electrode potential – electrode

equilibrium at p = 1 atm, t = 25 °C and Sia = 1

R is the universal gas constant (J/mol K)

F is the Faraday constant

z is the number of electrons involved in given

electrochemical process.

b) Redox equilibrium

Ox + ze- Red

Mz+ + ze- M (Cu2+ +2e- Cu)

ε = ε0 + +zMazFRT ln [208]

Or

0Cu/CuCu/Cu 22 εε ++ = + +2Culn a

zFRT [209]

By generalization:

0Cu/CuCu/Cu 22 εε ++ = +

red

ox

aa

zFRT ln [210]


4.3.3. Types of electrodes

Reversible electrodes are divided in three types: 1st type electrodes, 2nd

type electrodes and 3rd type electrodes according to the metal and electrolyte

nature.

a) 1st type electrode

a1) The electrode Metal – Ion (metal is introduced in its solution)

M Mz+ +ze- (heterogenous reaction)

0M/MM/M zz εε ++ = + +za

zFRT

Mln [211]

Examples: Cu/CuSO4, Zn/ZnSO4, Ni/NiSO4, Ag/AgNO3

a2) Gas electrode consists in an inert metal (platinum) immersed in a

solution which contains a gas dissolved (i.e. H2, O2 Cl2) at a given partial pressure

and ion of this gas (H+, HO-, Cl-) at a given concentration. The gas is partially

adsorbed on the platinum surface and establishes with its ions from solution a

chemical equilibrium. The Nernst expression for the electrode potential is given by:

0εε = + red

ox

aa

zFRT ln [212]

a21) Hydrogen electrode: (Pt) H2/H+ or (Pt) H2/HCl is characterized

by the following reactions:

H2(aq) 2H 2H+(aq) +2e-

0H/HH/H 22

εε ++ = + 2H

2Hln

2 pa

FRT + [213]

or 0H/HH/H 22

εε ++ = + 2/1H

H

2

lnpa

FRT + [214]

By convention for standard conditions (2Hp = 1atm, t = 25 °C) and

+Ha = 1 0

H/H 2ε + = 0

Electrochemistry 101

⇒ =+2H/H

ε 2/1H

H

2

lnpa

FRT + [215]

=+2H/H

ε – 0.059 pH (for ≠+Ha 1; t = 25 °C and

2Hp = 1atm)

a22) Oxygen electrode: (Pt) O2/HO- or (Pt) O2/NaOH is formed

from oxygen introduced under pressure in an electrolyte which contains hydroxide

ions, HO-, (i.e. NaOH) and a platinum bare:

O2 + 2H2O + 4e- 4HO-

0HO/OHO/O -

2-

2εε = + 4

HO

O2ln4 −a

pF

RT [216]

0HO/OHO/O -

2-

2εε = +

−HO

4/1O2ln

ap

FRT [217]

For 2Op = 1 atm and t = 25 °C

0HO/OHO/O -

2-

2εε = + −HO

lg059.0 a

a23) Chlorine electrode (Pt) Cl2/Cl-:

Cl2 + 2e- 2Cl-

0C/ClC/C 2

-2

εεlll

=− + 2C

Cl2ln2 −l

ap

FRT [218]

b) 2nd type electrodes have the general form:

M/MX, X-

Such types of electrodes consist of a metal (Ag, Hg) in direct contact

with an insoluble salt (or hardly soluble) of the above mentioned metal (AgCl,

Hg2Cl2). The whole system being immersed into a solution containing a soluble salt

of the anionic species involved (KCl).

The best known examples are the calomel electrode Hg/Hg2Cl2, KCl

and silver chloride electrode Ag/AgCl, KCl.

For the silver chloride electrode equilibrium reactions at the interface

are:


Ag Ag+ + e-

Ag+ + Cl- AgCl

∑: Ag + Cl- AgCl + e-

The electrode potential for this 2nd type of electrode could be obtained

starting from the potential expression for a 1st degree electrode, considering for

instance the silver chloride electrode as 1st degree electrode: Ag/Ag+. The silver

ions activity is deduced from the product of solubility of AgCl:

PS = −+ ⋅ ClAgaa [219]

−

+ =Cl

Ag aa SP [220]

⇒ 0/A,C/ εε +=

AggKCllAgAg + +Aga

FRT ln [221]

0/A,C/ εε +=

AggKCllAgAg + −Cl

aFRT SPln [222]

c) Redox electrodes consist in an inert metal immersed in a solution

which contains two substances able to pass from one in another by electronic

transfer. This electronic transfer is realized through the inert metal.

Examples:

Pt/Fe2+, Fe3+; Pt/Sn4+, Sn2+.

For the electrode Pt/Fe2+, Fe3+ the electrode reaction is:

Fe3+ +e- Fe2+

and the electrode potential is:

0// 2323 εε ++++ =FeFeFeFe

+ +

+

2

3

Fe

Flnaa

FRT e

Depending on the nature of ion which involved in the charge transfer

through the double electric layer there are:

− Reversible electrodes with respect to the cation (ex. 1st type

electrodes);


− Reversible electrodes with respect to the anion (ex. 2nd type of

electrodes);

- Redox electrodes having only electronic transfer to the double

layer.

4.4. Electrochemical cells

4.4.1 Definition. Redox reactions. Electromotive force. Electrode

Potential measurement.

An electrochemical (galvanic) cell is a device used for creating an

electromotive force (voltage) and the electrical current from chemical energy of

the oxido – reduction reactions.

The electrochemical cells are sources of direct current generated by the

redox reactions which take place spontaneously. The word redox summarized the

electrochemical reactions (reduction and oxidation)

The reverse process is the electrolysis, in which the chemical reaction takes

place due to the passing of the electric current. Table 2 Comparison between electrochemical (galvanic) cells and electrolytic cells

Electrochemical cells Electrolytic cells

• Transform chemical energy in

electrical energy

• Oxidation reaction takes place

at the anode which is symbolised

by negative sign (-)

• Reduction reaction takes place

at the cathode which is

symbolised by positive sign (+)

• Transform electrical energy in

chemical energy

• Oxidation reaction takes place at

the anode which is symbolised by

positive sign (+)

• Reduction reaction takes place at

the cathode which is symbolised

by negative sign (-)


Any oxidation reaction takes always place in tandem with a reduction

reaction. By oxidation a substance looses electrons while by reduction a substance

gains electrons. The number of lost electrons is always equal to the number of

gained electrons. Per total, the process is neutral from the electrical point of view.

Example:

Zn Zn2+ + 2e- – oxidation reaction which takes place at the anode

Cu2+ + 2e- Cu – reduction reaction which takes place at the cathode

∑: Zn + Cu2+ Zn2+ + Cu – the redox process (overall reaction)

A galvanic cell consists in:

• two compartments or two half – cells, each composed of an electrode

dipped in a solution of electrolyte. These half – cells are designed to

contain the oxidation and reduction half reactions;

• a porous barrier (salt bridge) preventing the spontaneous mixing of the

aqueous solution in each compartment, but allowing the migration of

ions in both directions in order to maintain electrical neutrality;

The two half-cells are also connected externally. In this way, electrons

provided by the oxidation reaction are forced to travel via an external circuit to the

site of the reduction reaction.

A galvanic cell is symbolized in the following way:

(–) M1/electrolyte1, +zMa1

// electrolyte2, +zMa2

/ M2 (+)

The fact that the reaction occurs spontaneously once these half cells are

connected indicates that there is a difference in potential energy between them.

This difference in potential energy is called electromotive force, emf, and is

measured in Volts.

Conventionally, the electromotive force is the difference between the

equilibrium potential of the positive electrode and the equilibrium potential of the

negative electrode.

E = εe+ - εe- [223]


The electrode potential is determined by using commonly a reference

electrode. An electrochemical cell with a common reference electrode and the

electrode with unknown electrode potential is formed and the emf is measured.

Knowing the electrode potential of the reference electrode one can determins the

electrode potential for the metal.

Common reference electrode:

− Hydrogen electrode ( 0H/H 2

ε + = 0 V)

− Mercury chloride (calomel) electrode ( 0εcal = 0.242 V)

4.4.2. The Volta electrochemical series

Alessandro Volta arranged the metals in a series after their ability to form

ions in solution taking hydrogen as reference. In this way the electrochemical

series were obtained taking advantage of the fact that the metals are also arranged

in order of their standard electrode potentials (redox potentials). The most negative

ε° values are placed on the left side of the electrochemical series, and the most

positive on the right side, respectively:

Li … Mg … Al … Zn … Fe … Ni … H … Cu ... Ag ... Au

(-3.01) (-2.38) (-1.66) (-0.76) (-0.44) (-0.23) (0.00) (+0.34) (+0.79) (+1.42) ε°, V

The electrochemical potential series has a special importance. After its

position in this series one can predict the chemical behaviour of metal in different

media (acid, basic or neutral).

Every metal replaces in its solutions the metals located on its right in the

electrochemical potential series.

Example: Al + CuSO4 Al2(SO4)3+Cu

The metals which are on the left side of hydrogen replace it in water and in

diluted acidic solutions while the metals placed on the right side do not.

Example 2Na + 2H2O 2NaOH + H2


Generally the metals placed on the left side of hydrogen are oxidized

relatively easy, react with acids and sometimes with water generating stable

combinations.

Theoretically the replacement of the hydrogen from water at pH = 7 by the

metals is a spontaneous reaction with release of hydrogen for the metals with

electrode potential less than – 0.414 V.

2/10

//2

22ln

H

HHHHH p

aF

RT +

++ += εε ; 0/ 2HH +ε = 0,

2Hp = 1atm

⇒ 2/ HH +ε =0,059× (– 7) = – 0,414 V

The more a metal is on the left, the more active it is, and the more difficult

is to reduce it. This means its ionization capacity is larger.

The metals situated on the right side of the hydrogen, called also noble

metals, present a reduced chemical activity. They oxidize with difficulty and they

have a big tendency to pass from ions state in atomic state.

Example Cu2+ + 2e- Cu

The electrochemical potential series reveal also that the electrochemical

potential depends on chemical nature of the metal and on its valence.

Example: 0Cu/Cu 2ε + = + 0.344 V; 0

Cu/Cuε + = 0.521 V

4.4.3. Reversible and irreversible galvanic cells

a) Reversible galvanic cell – Daniel Jacobi cell has the following

structure:

(–) Zn/ZnSO4//CuSO4/Cu (+)

With the following half reactions:

Anode (–) oxidation reaction: Zn Zn2+ + 2e-

Cathode (+) reduction reaction: Cu2+ + 2e- Cu

∑: Zn + Cu2+ Zn2+ + Cu

or Zn + CuSO4 ZnSO4 + Cu


If to this device we apply a larger voltage, with opposite sign than emf then

we will transform the electrochemical cell into an electrolytic cell. In such cell the

reverse reactions will take place:

- copper will be dissolved in CuSO4 solution

Cu Cu2+ + 2e-

- zinc will be deposed on Zn electrode

Zn2+ + 2e- Zn

Conditions for a reverse cell:

- when the circuit is open, no chemical reaction will occur;

- cell reaction which takes place when a little bit higher external voltage

with respect to emf is applied should be the reverse of that which

appears when the cell produce voltage;

- during the function the current should be not very high and the emf

should be constant.

b) Irreversible galvanic cell – Volta cell has the following structure:

(–) Zn/H2SO4/Cu (+)

with the following half reactions:

Anode (–) oxidation reaction: Zn Zn2+ + 2e-

Cathode (+) reduction reaction: 2H+ + 2e- H2 ↑

∑: Zn + 2 H+ Zn2+ + H2 ↑

or Zn + H2SO4 ZnSO4 + H2 ↑

It can be easily observed that the current will have the direction from

anode to cathode.

If we apply an external voltage, a little bit higher then emf and with

opposite sign the copper electrode will be dissolved (Cu Cu2+ + 2e-) and

the hydrogen will be reduced on the zinc electrode (2H+ + 2e- H2 ↑). So

the reverse reactions are not produced.

An electrochemical cell is irreversible if:


- there are chemical transformations even when the cell does not

produce current in open circuit;

- the half reactions which occurs when the electrochemical cell is

transformed in electrolytic cell is not the reverse of that taking place

when the cell produces the current;

- Electromotive force is decreasing quickly when cell produces current.

4.5. Applications of electrochemistry in electronics

Driven by the electronics industry, electrochemical technology has rapidly

evolved, finding increasing applications in microelectronics, batteries, sensors,

materials science, industrial fabrication, corrosion, microbiology, neurobiology and

medicine.

Semiconductors are probably the most used materials in electronics.

Electrochemical processes in semiconductors are practiced today in the fabrication

of microelectronics products. Electrochemical technology has been an integral part

of major advances in semiconductor technology of the 20th century. While, the

microelectronics industry is rapidly reaching maturity, its explosive start and

sustained growth has been partly enabled by the use of scalable electrochemical

processes.

Prior to the invention of the transistor and the birth of the semiconductor

industry, the field of electrochemistry was already very advanced with respect to

both theoretical understanding and industrial applications. Most semiconductors

materials for opto-electronic applications must be in the form of single crystals

with exceptional crystalline perfection and purity. Some semiconductors in

polycrystalline film or bulk form have also been found useful in a few applications,

the most important being low cost solar cells. This latter application has stimulated

much of the recent work on the electrolytic deposition of semiconductors materials.


Traditional macroscopic electrochemistry started 200 years ago, presenting

technological connection to energy technology and electrolysis. In the last 10

years, electrochemical nanotechnology has been favored and it is, in fact, a field for

fundamental science. As a link between these different fields, micro

electrochemistry is a fast growing part of electrochemistry, forming

interdisciplinary bridges to science and medicine. Figure 26 shows the

interdisciplinary character of electrochemical microsystems. These systems include

and develop the fundamental knowledge of electrochemistry, but in difference to

fundamental science.

Figure 26 Electrochemical microsystem technologies to five special topics of fundamentals,

galvanics, systems, analysis and biology. Many special subjects are located in between, because they belong to various topics (J.W. Schultze)

System means a combination of two or more parts or devices. A single

microelectrode is essential for an electrochemical microsystem, but it represents

only a part of it. Due to the large applications, electrochemical technologies include

electrochemical processes. When we talk about electrochemical system in

electronics, the scale ranges to micro domain. Micro limits the systems and

processes to the range below 1 mm in all three dimensions: x, y, z. This yields a


clear difference to the well known field of thin film technology which limits only

the vertical dimension (z). For example, the miniaturization to ultra microelectrode

of nm dimensions is in progress with common technologies.

Microgalvanics – The technologies of microgalvanics are now well

established. Electrochemical Microsystems started in electronics with the invention

of the printed circuit board in the mm-range, but the dimensions were continuously

decreased due to the requirements of electronics. The dominating technology is

now lithography. The number of interconnections increased, and they had to be

miniaturized. While galvanic processes are applied in the nm-range for micro and

nano electronics, packaging in µm dimensions is still a limiting factor.

Simultaneously, the development of the new processes opened up the

electrochemical mass production of micro mechanic and opto-electronic tools,

which now enhance the development of micro reactors for chemical synthesis,

mixers and other tools.

Microelectrodes and microcells

During the last year, large progress were unregistered in the process of

miniaturization of electrodes, that was not done without some problems coming from

mechanism of electrode production, sensitivity and insulation or leakage currents.

Microelectrodes successfully found their appliance in electroanalysis and studies of

nucleation and growth. Even in biological research, pH sensitive electrodes in glass

technology have been applied. The resolution measured by the tip is in atomic

dimensions, while the exposed electrode surface remains in the µm range.

Another interesting direction in electrochemistry applied in the field of

electronics was the developing of small instrumentation market and software

intended for military and consumer applications. The only big time product

remains the pH-meter. To predict the future it is enough to take a look to

telecommunications, including the integration of video with computers, worldwide

wireless communications, advanced software development tools, and further

miniaturization of hardware by great integration of digital and analogue modules.

Even now we can use processor-based instrumentation to carry out an array of


electrochemical techniques on a given sample. One of the most productive aspects

of this approach is the fact that multiple users can employ an electrochemical

workstation without interfering with each other.


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Curs Chimie Cap4