1. Electrochemistry Prof. Dr. Sabine Prys @designed by ps
@designed by ps

2. Contents Chemical Reactions Chemical Reaction Kinetics
Chemical Reaction Thermodynamics
Chemical Solutions
Electrochemical Reactions
Galvanic Cells & Batteries
Standard Potential
Specific Electrodes
NERNST’s equation
Corrosion

3. 1.0 Chemical Reactions redox reaction Anodic oxidation Ag  Ag+ + e-
Fe2+  Fe3+ + e-
Cathodic reduction
Ag+ + e-  Ag
Fe3+ + e-  Fe2+
stochiometrical yields
Law of mass conservation !

4. Driving Forces
Kinetics (chemical process rates) Thermodynamics (chemical process energy flow) reaction energy change: standard enthalpy* DH298 DH = DU + p.DV + vDp DU = DQ + DW chemical reactions in open vessels: DH ≈ DU * enthalpy at standard conditions: T = 298 K, p= 1,013 bar
Internal energy of a system plus the product of pressure and volume. Its change in a system is equal to the heat brought to the system at constant pressure.
Quantity the change in which is equal to the sum of heat Q brought to the system and work W performed on it DU = Q +W . Also calledthermodynamic energy.
H enthalpy U internal energy Q flow of heat W Work
T temperature P pressure V volume

5. 2.0 Chemical Reaction Kinetics
reversible chemical reaction e.g. in aqueous solution
[A],[B] = reactant concentration a,b = mols of corresponding reactants
[C],[D] = product concentration c,d = mols of corresponding products
k = velocity of reaction k = velocity of back reaction
K = equilibrium constant
equilibrium constant
Quantity characterizing the equilibrium of a chemical reaction and defined by an expression of the type , where  is the stoichiometric number of a reactant (negative) or product (positive) for the reaction and  stands for a quantity which can be the equilibrium value either of pressure, fugacity, amount concentration, amount fraction, molality, relative activity or reciprocal absolute activity defining the pressure based, fugacity based, concentration based, amount fraction based, molality based, relative activity based or standard equilibrium constant (then denoted), respectively.

6. (reversible) chemical reaction
2.1 Equilibrium Constant
(reversible) chemical reaction
Kx = equilibrium constant
xB = equilibrium value either of pressure, fugacity, amount concentration, amount fraction, molality, relative activity or reciprocal absolute activity defining the pressure based, fugacity based, concentration based, amount fraction based, molality based, relative activity based or standard equilibrium constant
B = stoichiometric number of a reactant (negative) or product (positive)
fugacity, , 
Of a substance B  or in a gaseous mixture is defined by ,
F = f*p
where  p is the partial pressure of B and f its fugacity coefficient
f depends on how much the real gas behaviour deviates from the ideal gas behavior.

7. 2.1.1 Reaction Velocity 1st order kinetics velocity of reaction [c]

8. Mass-Law Effect
chemical potential of reaction () = chemical potential of back reaction ()
equilibrium constant K depends on temperature
chemical equilibrium: concentrations are not independent
The law stating that the rate of any given chemical reaction is proportional to the product of the activities (or concentrations) of the reactants.

9. 2.1.3 Example: Neutralisation Reaction
Acid Base Reaction :
acis + base = salt + water
But also: +
Säure-Neutralisation
Gibt man zu Salzsäure die gleiche Stoffmenge an Natronlauge, so entsteht eine neutrale Kochsalzlösung:
 NaOH  +  HCl       NaClaq  +  H2O
 NaCl aq    Na+ aq + Cl- aq
Aus den Hydronium- und Hydroxid-Ionen hat sich durch Protonenübergang Wasser gebildet, während die Na+- und Cl--Ionen unverändert vorliegen. Die entsprechende Lösung reagiert dann neutral. Die Aufhebung von Säureeigenschaften durch eine Base bzw. umgekehrt wird allgemein als Neutralisation bezeichnet. Bei der Reaktion von Säure mit Base entstehen immer die Reaktionsprodukte Wasser und Salz.

10. 3.0 Chemical Reaction Thermodynamics
one component thermodynamic system: e.g. Ideal gas
U Internal energy – total energy contained by a one component system
Quantity the change in which is equal to the sum of heat Q brought to the system and work W performed on it DU = Q +W . Also calledthermodynamic energy.
U internal energy Ekin kinetic energy Epot potential energy
Q flow of heat W Work T temperature
S entropy p pressure V volume

11. 3.1 „Ideal“ Gas
one component thermodynamic system: e.g. Ideal gas
e.g. N2, H2, O2, He, Ne,....
V large, p small
System contains only one particle type
System contains statistically sufficient particles
System contains point particles
System contains non interacting particles
System contains randomly moving particles
All particle collisions are elastic

12. 3.1 „Ideal“ Gas Law one component thermodynamic system: e.g. gas
p = pressure [Pa] = [N/m2]
V = volume [m3]
n = number of mols [mol-1]
R = gas constant [J/(mol+K)]
T = Kelvin temperatur [K]
gas constant: R = 8,3145 J.K-1.mol-1

13. 3.2 The BOLTZMANN Constant
R = gas constant NA = Avogadro constant
The BOLTZMANN constant  is the physical constant relating energy at the individual particle level with temperature observed at the collective or bulk level

14. 3.3 Normal & Standard Conditions
Such definitions can vary according to different sources: IUPAC, NIST, …
normal conditions: normal pressure p = 1 atm = 101,325 kPa = 1013,25 mbar normal temperature T = 0°C = K
 standard conditions:
standard pressure p = 1 atm = 101,325 kPa = 1013,25 mbar standard temperature T = 25°C = K

15. 3.4 Enthalpy
Enthalpy is a measure of the total energy of a thermodynamic system including
the internal energy (energy required to create a system),
the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure.
Enthalpy is a thermodynamic potential, a state function and an extensive quantity (i.e. depending on amount material).
In the physical sciences, an intensive property (also called a bulk property, intensive quantity, or intensive variable), is a physical property of a system that does not depend on the system size or the amount of material in the system: it is scale invariant.
By contrast, an extensive property (also extensive quantity, extensive variable, or extensive parameter) of a system is directly proportional to the system size or the amount of material in the system.
For example, density is an intensive property of a substance because it does not depend on the amount of that substance; mass and volume, which are measures of the amount of the substance, are extensive properties.
Note that the ratio of two extensive properties that scale in the same way is scale-invariant, and hence an intensive property.
wikipedia
H enthalpy U internal energy
P pressure V volume

16. Exercise: gas volume
Calculate the volume of 1 mol of an ideal gas !
under normal conditions
under standard conditions

17. 3.4.2 Exercise: internal energy
What is the internal energy for 1 Mol He at 20°C ?

18. 3.4.3 Exercise: internal energy change
What is the internal energy change for 1 Mol He at 20 ° C , p = 1000 Pa and DQ = 0 when expanding the volume from 1 m3 to 2 m3 ?

19. Exercise: enthalpy
What is the enthalpy for 1 Mol He at 20 ° and standard pressure ?

20. 3.5 Entropy
Quantity the change in which is equal to the heat brought to the system in a reversible process at constant temperature divided by that temperature.
Entropy is zero for an ideally ordered crystal at 0 K . 
In statistical thermodynamics where k  is the Boltzmann constant and W  the number of possible arrangements of the system the entropy is
DS = k . lnW
Entropy is a thermodynamic property that is a measure of the energy not available for work in a thermodynamic process. It is defined by the second law of thermodynamics. In the microscopic interpretation of statistical mechanics, entropy expresses the disorder or randomness of the constituents of a thermodynamic system or, analogously, the availability of accessible quantum mechanical states. A closed system always tends towards achieving a state with a maximum of entropy.

21. 3.5.1 Exercise: entropy change none
What is the entropy change for 1 Mol He at 20 ° C , p = 1000 Pa and DQ = 0 when expanding the volume from 1 m3 to 2 m3 ?

22. 3.5.1 Exercise: entropy change
What is the entropy change for 1 Mol He at 20 ° C when the flow of heat DQ = J ?

23. 3.6 Chemical Reaction Energy
multi component thermodynamic system: e.g. Chemical reaction
Internal energy of a one component thermodynamic system
Internal energy of multi-component systems
U internal energy
Ekin kinetic energy
Epot potential energy
Q flow of heat
W Work
T temperature
S entropy
P pressure
V volume
Ni component i
µi chemical potential
G Gibbs’ energy
Quantity the change in which is equal to the sum of heat Q brought to the system and work W performed on it DU = Q +W . Also called thermodynamic energy.

24. 3.7 Enthalpies of Multi Component Systems
Enthalpy of reaction
Enthalpy of formation
Enthalpy of combustion
Enthalpy of hydrogenation
Enthalpy of atomization
Enthalpy of neutralization
Enthalpy of solution
Enthalpy of hydration
Enthalpy of fusion
Enthalpy of vaporization
Enthalpy of sublimation
Lattice enthalpy
…..

25. Endothermic Reaction
Endothermic A reaction for which the overall standard enthalpy change DH298 >0 e.g. t E
activation energy
+ DH

26. Exothermic Reaction
Exothermic A reaction for which the overall standard enthalpy change DH298 < 0 e.g. E
activation energy
- DH t

27. 3.8 GIBBS’ Free Energy G
Useful energy, or energy available to do work G
G = free energy H = GIBBS’ energy (enthalpy)
U = internal energy T = Kelvin temperature
S = entropy p = pressure
V = volume
TDS is the energy not available for doing work

28. 3.9 Spontaneity of Redox Reactions
DH DS Spontaneity
Exothermic DH < 0 Increase DS > DG < 0
Exothermic DH < 0 Decrease DS < if |TDS| < |DH|
Endothermic DH > 0 Increase DS > if TDS > DH
Endothermic DH > 0 Decrease DS < DG > 0

29. 3.10 Thermodynamical Equilibrium
Reversible processes ultimately reach a point where the rates in both directions are identical, so that the system gives the appearance of having a static composition at which the Gibbs energy G is a minimum
DG = 0
At equilibrium the sum of the chemical potentials of the reactants equals that of the products, so that:
DG = DG298 + RT . lnK = 0  DG298 = - RT . lnK
The equilibrium constant K is given by the mass-law effect.

30. 3.10.1 Examples 6 C (s) + 3 H2 (g)  C6H6 (l) DG298 = + 124 kJ
C6H6 (l)  6 C (s) H2 (g) DG298 = kJ
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O DG298 = - 959,42 kJ

31. reduction = electron acceptance
Reduction of SiO2
reduction = electron acceptance
production of metallurgical grade silicon Electric arc furnace with carbon electrodes 14 kWh / kg Si !!!

32. oxidation = electron donation
SiHCl3 Synthesis
oxidation = electron donation
oxidation of silicon to SiHCl3
Thermische Oxidation
Bei der thermischen Oxidation werden die Siliciumwafer bei ca. 1000 °C in einem Oxidationsofen oxidiert. Dieser Ofen besteht im Wesentlichen aus einem Quarzrohr in dem sich die Wafer auf einem Carrier aus Quarzglas befinden, drei bis fünf getrennt regelbaren Heizwicklungen und verschiedenen Gaszuleitungen. Das Quarzglas hat einen sehr hohen Schmelzpunkt (weit über 1500 °C) und ist deshalb sehr gut für Hochtemperaturprozesse geeignet. Damit es nicht zu Scheibenverzug oder Scheibensprüngen kommt, wird das Quarzrohr in sehr kleinen Schritten (max. 10 °C pro Minute) aufgeheizt. Die Temperierung ist mittels der getrennten Heizwicklungen im gesamten Rohr über eine Länge von ca. 1 m auf ± 0,5 °C exakt regelbar.
Darstellung eines Oxidationsofens
Der Sauerstoff strömt dann als Gas über die Wafer und reagiert an der Oberfläche zu Siliciumdioxid. Es entsteht eine glasartige Schicht mit amorpher Struktur. Je nach Prozessgas finden dann verschiedene Oxidationen statt (eine thermische Oxidation muss naturgemäß auf einer Siliciumoberfläche stattfinden). Die thermische Oxidation unterteilt sich in die trockene und feuchte Oxidation, welche sich wiederum in die nasse Oxidation und die H2O2-Verbrennung gliedern lässt.
Trockene Oxidation:
Der Oxidationsprozess findet unter reiner Sauerstoffatmosphäre statt. Dabei reagiert Silicium mit Oxid zu Siliciumdioxid: Si + O2→SiO2
Dieser Prozess findet in der Regel bei  °C statt. Zur Erzeugung von sehr stabilen und dünnen Oxiden wird die Oxidation bei ca. 800 °C durchgeführt.
Eigenschaften der Trockenoxidation:
langsames Oxidwachstum / hohe Dichte / hohe Durchbruchspannung (Verwendung von elektrisch stark beanspruchten Oxiden, z.B. Gateoxid)
Nasse Oxidation:
Bei der nassen Oxidation wird der Sauerstoff durch ein Bubbler-Gefäß mit Wasser (ca. 95 °C) geleitet, so dass sich zusätzlich zum Sauerstoff auch Wasser in Form von Wasserdampf im Quarzrohr befindet. Daraus ergibt sich folgende Reaktionsgleichung: Si + 2H2O→SiO2 + 2H2
Dieser Prozess findet bei  °C statt. Die nasse Oxidation weist folgende Eigenschaften auf: hohes Oxidwachstum schon bei geringer Temperatur (Verwendung von Maskierschichten, Feldoxid) geringere Qualität als Trockenoxid
H2O2-Verbrennung:
Bei der H2O2-Verbrennung wird neben hochreinem Sauerstoff auch hochreiner Wasserstoff verwendet. Die beiden Gase werden getrennt in das Quarzrohr geleitet und an der Eintrittsöffnung verbrannt. Damit es nicht zu einer Knallgasreaktion mit dem hochbrennbaren Wasserstoff kommt, muss die Temperatur über 500 °C liegen, die Gase reagieren dann in einer stillen Verbrennung. Dieses Verfahren ermöglicht die Erzeugung von schnell wachsenden und nur wenig verunreinigten Oxidschichten. Damit lassen sich sowohl dicke Oxide, als auch dünne Schichten bei vergleichsweise geringer Temperatur (900 °C) herstellen. Die niedrige Temperatur erlaubt auch die thermische Belastung von bereits dotierten Wafern (siehe Dotieren mittels Diffusion).Bei allen thermischen Oxidationen ist das Oxidwachstum auf 111-orientierten Substraten höher als auf 100-orientierten (siehe der Einkristall). Außerdem erhöht ein sehr hoher Anteil an Dotierstoffen im Substrat das Wachstum deutlich.
Ablauf des Oxidationsvorgangs:
Zu Beginn reagiert der Sauerstoff an der Waferoberfläche zu Siliciumdioxid. Nun befindet sich eine Oxidschicht auf dem Substrat durch die der Sauerstoff zunächst diffundieren muss, um mit dem Silicium reagieren zu können. Die Aufwachsrate hängt nur zu Beginn von der Reaktionszeit zwischen Silicium und Oxid ab; ab einer gewissen Dicke wird die Oxidationsgeschwindigkeit von der Diffusionsgeschwindigkeit des Oxids durch das Siliciumdioxid bestimmt. Mit zunehmender Oxiddicke verlangsamt sich also das Wachstum. Da die entstandene Schicht amorph ist, sind nicht alle Bindungen der Siliciumatome intakt; es gibt teilweise freie Bindungen (freie Elektronen und Löcher) an der Si-SiO2-Grenzschicht. So ergibt sich an diesem Übergang insgesamt eine leicht positive Ladung. Da sich diese Ladung negativ auf Bauteile auswirken kann wird versucht sie so gering wie möglich zu halten. Das kann beispielsweise mit höherer Oxidationstemperatur erreicht werden, oder durch Verwendung der nassen Oxidation, die ebenfalls nur eine sehr geringe Ladung verursacht.
Segregation:
Bei der thermischen Oxidation wird Silicium, durch die Reaktion mit Sauerstoff zu Siliciumdioxid, verbraucht. Das Verhältnis der aufgewachsenen Oxidschicht zu verbrauchtem Silicium beträgt 2,27; d.h. das Oxid wächst zu 45 % der Oxiddicke in das Substrat ein.
Aufwachsverhalten von Oxid auf Silicium
Dotierstoffe die sich im Substrat befinden, können im Siliciumkristall oder im Oxid eingebaut werden, dies hängt davon ab, in welchem Material sich der Dotierstoff besser löst. Das Verhalten lässt sich mit einer Formel berechnen, k wird dabei als Segregationskoeffizient bezeichnet:
Ist k größer 1 werden die Dotierstoffe an der Oberfläche des Substrats eingebaut, bei k kleiner 1 reichern sich die Dotierstoffe im Oxid an.
Oxidation durch Abscheidung
Bei der thermischen Oxidation wird Silicium des Wafers zur Oxidbildung verbraucht. Ist die Siliciumoberfläche jedoch durch andere Schichten verdeckt, muss man das Oxid über Abscheideverfahren aufbringen, bei denen neben Sauerstoff auch Silicium selbst hinzugefügt wird. Die zwei wichtigsten Verfahren dabei sind die Silanpyrolyse und die TEOS-Abscheidung. Eine ausführliche Beschreibung der Verfahren folgt im Kapitel Abscheidung.
Silanpyrolyse:
Pyrolyse bedeutet, dass chemische Verbindungen durch Wärme gespaltet werden, in diesem Fall das Gas Silan SiH4 und hochreiner Sauerstoff O2. Da sich das giftige Silan bei einer Konzentration von 3 % in der Umgebungsluft selbst entzündet muss es mit Stickstoff oder Argon auf 2 % verdünnt werden. Bei ca. 400 °C reagiert das Silan und der Sauerstoff zu Siliciumdioxid und Wasserstoff:SiH4 + O2→SiO2 + 2H2
Das Siliciumdioxid ist nur von geringer Qualität. Alternativ kann eine Hochfrequenzanregung über ein Plasma bei ca. 300 °C zur Dioxidabscheidung verwendet werden. So entsteht ein etwas stabileres Oxid.
TEOS-Abscheidung:
Das bei dieser Methode verwendete Tetraethylorthosilicat, kurz TEOS (SiO4C8H20) enthält die beiden benötigten Elemente Silicium und Sauerstoff. Unter Vakuum geht die bei Raumtemperatur flüssige Verbindung bereits in Gasform über. Das Gas wird in ein beheiztes Quarzrohr überführt und dort bei ca. 750 °C gespalten.
SiO4C8H20→SiO2 + Nebenprodukte Das Siliciumdioxid scheidet sich auf den Wafern ab, die Nebenprodukte (z.B. H2O - Wasserdampf) werden abgesaugt. Die Gleichmäßigkeit dieses Oxids wird durch den Druck im Quarzrohr und die Prozesstemperatur bestimmt. Es ist elektrisch stabil und sehr rein

33. 3.11 Chemical Potential 1 G = Gibbs free energy n+ = moles of cations
nb = moles of type b n- = moles of anions
nj = moles of type j µb = chemical potential of type b
T = Kelvin temperature µ+ = chemical potential of cations
p = pressure µ- = chemical potential of anions

34. 3.11.1 Standard Chemical Potential
R = gas constant
T = Kelvin temperature
a = activity of an ion type
chemical potential µi:
change in the energy of the system when an additional constituent particle i is introduced, with the entropy and volume held fixed
Ausarbeiten !

35. 3.11.2 Chemical Potential 2 www.goldbook.iupac.org Gibbs energy:
Enthalpy minus the product of thermodynamic temperature and entropy. It was formerly called free energy or free enthalpy.

36. 3.12 Maximum Work Wmax = maximum work G = Gibbs free energy
R = gas constant
T = Kelvin temperature
K = equilibrium constant
z = ion charge
n = moles
F = Faraday‘s constant
E = galvanic cell potential
U = voltage
I = currant
t = time

37. 3.12.1 Examples What is the vaporisation enthalpy of water at 298 K ?
What is the heating energy of 1 m3 methane ?
What is the combustion energy of 1 m3 methane ?
Vaporisation enthalpy= (890 – 802)/2 = 44 kJ/mol
Heating energy = 1000 * 802 / 22,4 = kJ (273,15)
Combustion energy = 1000*890/22,4 = kJ (273,15)

38. 4.0 Chemical Solutions Suspension (particle diameters 10-4 - 10-5 cm )
solid particles in homogeneous fluid
Colloid (particle diameters cm )
microscopically dispersed particles in another substance
Solution (particle diameters cm )
Homogeneous phase with at least 2 components:
solvent and solute
gas in liquid e.g. O2 in H2O
gas in solid e.g. H2 in palladium
liquid in liquid e.g. petroleum
solid in liquid e.g. NaCl in H2O
 electrolytes in water

39. 4.1 Solvation & Dissolution
IUPAC: any stabilizing interaction of a solute … and the solvent … Such interactions generally involve electrostatic forces and van-der-Waals forces, as well as chemically more specific effects such as hydrogen bond formation.
Dissolution
IUPAC: mixing of two phases with the formation of one new homogeneous phase (i.e. the solution)
kinetic process, quantified by its rate
Any stabilizing interaction of a solute (or solute moiety) and the solvent or a similar interaction of solvent with groups of an insoluble material (i.e. the ionic groups of an ion-exchange resin). Such interactions generally involve electrostatic forces and van der Waals forces, as well as chemically more specific effects such as hydrogen bond formation.

40. 4.2 Solubility
IUPAC:
The analytical composition of a saturated solution, expressed in terms of the proportion of a designated solute in a designated solvent, is the solubility of that solute. The solubility may be expressed as a concentration, molality, mole fraction, mole ratio, etc.
quantifies the dynamic equilibrium state achieved when the rate of dissolution equals the rate of precipitation
solubility of salts: lattice energy & solvation enthalpy
ideal dilute solution

41. Lattice Energy
Compound Lattice Energy [kJ/mol] LiF -1019,9 LiCl - 838,6 LiBr - 798,8 LiI - 742,3 NaCl - 766,2 KCl RbCl - 665,7 CsCl - 649,4
Compound Lattice Energy [kJ/mol] MgO 3931,4 CaO 3479,2 SrO 3207,1 BaO 3043,8 Al2O ,2 CaF2 2612,6 CaCl2 2147,8 CaBr2 2026,4 CaI2 1921,7
Elektrolytgleichgewichte; Ackermann et. Al

42. 4.3 Aqueous Solutions
Solvatation: lattice disintegration break of hydrogen bonds  „hole“ formation formation of ion dipole interactions
Compound DHsolution_∞ Compound DHsolution_∞ Compound DHsolution_∞
LiF ,6 NaF + 2, KF ,2
LiCl ,0 NaCl + 5, KCl + 18,4
LiBr ,5 NaBr + 0, KBr + 21,4
LiI ,0 NaI , KI ,4
DHsolution_∞ = solution enthalpy in very low concentrations
hydrogen bond
A form of association between an electronegative atom and a hydrogen atom attached to a second, relatively electronegative atom. It is best considered as an electrostatic interaction, heightened by the small size of hydrogen, which permits proximity of the interacting dipoles or charges. Both electronegative atoms are usually (but not necessarily) from the first row of the Periodic Table, i.e. N, O or F. Hydrogen bonds may be inter-molecular or intramolecular. With a few exceptions, usually involving fluorine, the associated energies are less than kJ mol −1 (5 - 6 kcal mol −1).
Elektrolytgleichgewichte; Ackermann et. Al

43. 4.4.1 Hydratisation Enthalpies at 25°C
Ion DHHydr Ion DHHydr Ion DHHydr [J/mol] [J/mol] [J/mol] H ,4 NH ,7 OH ,2 Li ,7 Mg ,3 F ,3 Na ,8 Ca ,5 Cl ,8 K ,9 Br ,3 Rb ,1 I ,6 Cs ,9 Increasing ion radius  less energy gain !
Elektrolytgleichgewichte; Ackermann et. Al

44. 4.5 Saturated Solutions
IUPAC: A solution which has the same concentration of a solute as one that is in equilibrium with undissolved solute at specified values of the temperature and pressure. loading capacity (maximum loading, saturation capacity saturation loading) The maximum concentration of solute(s) that a solvent can contain under specified conditions. - +
solute

45. Gas Solutions
IUPAC:
The partial pressure (fugacity) of a solute (B) in a solution is directly proportional to the rational chemical activity (ax) of the solute:
Henry's law:
pB = partial pressure ax,B = chemical activity, ax,B∞ = solubility coefficient

46. 4.5.2 Solubility Coefficients of Gases in H2O
Temp.[°C] O2 N2 CO2 [g(Gas)/kg(H2O)/bar] 0 0,0676 0,0281 3, ,0526 0,0226 2, ,0428 0,0190 1, ,0364 0,0166 1, ,0291 0,0137 0, ,0258 0,0129 0, ,0246 0,

47. 4.5.3 Solubility of Salts in H2O
In mass %
Compound 20 °C 80 °C NaCl 26,5 27,5 KCl 25,5 33,6 NH4Cl 27,0 40,0 K2SO4 10,0 17,5 CaSO4 0,199 0,10 Ca(OH)2 0,17 0,087 CaCO3 0,0015 0,002 (100 °C) ZnCl2 78,7 84,5
Solubility depends on: ion radii, solvent dielectricity constant

48. 4.5.4 Densities of NaCl Solutions at 20°C
Concentration [mass%] Density [kg/m³]

49. 4.5.6 Solubility Curves for Unhydrated Salts
Often heat production
when dissolving non hydrated salts
e.g CaCl2, Na2CO3
figure reference: Elektrolytgleichgewichte; Ackermann et. Al

50. 4.5.7 Solubility Curves for Hydrated Salts
Often heat consumation
when dissolving hydrated salts
e.g CaCl2 * 6 H20, Na2CO3 * 10 H2O
Change of solid structure &
composition
figure reference: Elektrolytgleichgewichte; Ackermann et. Al

51. 4.5.8 Solubility Equilibrium
IUPAC: solubility product The product of the ion activities raised to appropriate powers of an ionic solute in its saturated solution expressed with due reference to the dissociation equilibria involved and the ions present.

52. 4.5.9 Low & High Solubility at 25°C
Compound Ksol AgCl 1, mol2/l2 AgBr 5, mol2/l2 AgI 8, mol2/l2 Ag2CrO4 1, mol2/l2 Ag2S 5, mol2/l2 Fe(OH)2 1, mol2/l2 Fe(OH)3 6, mol2/l2 FeS 4, mol2/l2  low solubility
Compound Ksol NaHCO3 1, mol2/l2 MgCO3 1, mol2/l2 PbCl2 1, mol2/l2 BaF2 1, mol2/l2 Ba(OH)2 5, mol2/l2 CaSO4 2, mol2/l2 Ag2SO4 1, mol2/l2  high solubility

53. 4.6 Ion activity
High ion concentrations in aqueous solutions ð ion – ion interactions:
pH measured < pH calculated
(1m, 0.1 m solution of acids)
ion activity:
a = activity, f = activity coefficient, c = concentration
f (HCl, 25°C): 0.001m/ m/ m/ m/0.809
Werte überprüfen !!!

54. 4.7 Colloidal Solutions
Larger particles in solvent, e.g. macromolecules / polymers
Properties depend on solute size and not on solute concentration !
Coagulation: growth of larger particles by smaller particles consumption
Hydrophobe colloids: large surface, large adsorption properties
Hydrophile colloids

55. 4.8 Gels

56. 4.9 Electrolytes
Electrolyte: solution which conducts electrical current Hydrated H3O+ Hydrated OH-

57. 4.9.1 Electrical Conductivity in Solutions
Electrolytes
solutions which support ion transport
salts in aqueous solutions, e.g. KCl, ZnSO4, CuCl2, etc.
molten salts
Conductivity L (resistance R)
bad electrolyte: distilled water:
µS/cm at 25 °C.
cathode
cat ions
anode
anions _ + -
H2O
[S] = [1/Ohm] Siemens

58. 4.9.2 Specific Conductivity
absolute electrolyte conductivity
R = solution resistance
specific electrolyte conductivity
A = electrode surface, l = electrode distance

59. 4.9.3 Example: Proton Migration
Grotthuss Diffusion
structural defect migration
mesomeric structures between H9O4+ and H5O2+,
Protonenwanderung
Da die Diffusion von Protonen etwa 5 mal schneller abläuft als eine entsprechende Diffusionen von Natrium- oder Chloridionen, wird von einem grundsätzlich anderem Mechanismus ausgegangen. Bei dieser sogenannten Grotthuss-Diffusion wandern nicht die Moleküle, sondern ein Strukturdefekt, was man sich vereinfacht als ein "Umklappen von Wasserstoffbrücken" vorstellen kann. Eine Vorstellung über diese Protonenwanderung und die Struktur des Defekts konnte durch Computersimulationen am Max-Planck-Institut für Festkörperforschung im Jahr 2000 formuliert werden: demnach verändert der protonische Defekt während der Migration seine Struktur kontinuierlich. Die Strukturgrenzformen bewegen sich zwischen H9O4+ und H5O2+, mitunter erweist sich der delokalisierte Bereich aber auch über mehrere Wasserstoffbrücken verteilt. Damit konnte neues Licht auf eine 200 Jahre alte Diskussion über die "wahre Struktur des Wassers" geworfen werden.

60. 5.0 Electrochemical Cells
Cl2 H2 + -
electrolytic cell galvanic cell
2 HCl (aq) ð H2 (g) + Cl2 (g) H2 (g) + Cl2 (g) ð 2 HCl (aq)
electrical energy  chemical energy chemical energy  electrical energy

61. 5.1 Electrolysis
electrolysis: decomposing materials by electric current
H2SO4 + 2H2O ð 2H3O+ + SO42-
water electrolysis
cathodic reduction
4H3O+ + 4e- ð 2 H2 ñ + 2 H2O
anodic oxidation
4 OH- ð 2 H2O + O2 ñ + 4 e-
total
2 H2O (l) ð 2 H2 (g) + O2 (g)
H20 + H2SO4
1:10
electrods
battery
ca. 15 V

62. 5.1.1 Electrochemical Equivalent
Q = electric charge in C
n = yield in mol
F = Faraday‘s constant
= ,309 As / mol
Ec = electrochemical equivalent
M = ion weight
z = ion charge
NL = Lohschmidt‘s number
e = elementary charge

63. 5.1.2 Faraday‘s Laws m = Ec . Q = Ec .I. t m = mass yield in g
Ec = electrochemical equivalent
Q = electric charges in Coulomb
I = current strength
t = electrolysis time
ma ,mb = mass yield in g
for material a / b
Ma,Mb = molecular weight for material a / b
za, zb = chemical valency for material a / b

64. 5.2 Galvanic Elements Daniell Element: 2 galvanic half cells + bridge
voltmeter ca 1,1 V Zn Cu
1 m
ZnSO4
CuSO4
diaphragma
(pottery)
bridge containing
KCl solution
Daniell Element:
2 galvanic half cells + bridge
Zn / ZnSO4 // CuSO4 / Cu
electrode reactions
Zn (cathode) ð Zn2+ + 2e-
Cu2+ + 2e- ð Cu (anode)
Zn metal in ionic solution
Cu ions in Cu metal
electrical current results from different oxidation affinities
To provide a complete electric circuit, there must also be an ionic conduction path between the anode and cathode electrolytes in addition to the electron conduction path. The simplest ionic conduction path is to provide a liquid junction. To avoid mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while reducing electrolyte mixing. To further minimize mixing of the electrolytes, a salt bridge can be used which consists of an electrolyte saturated gel in an inverted U-tube. As the negatively charged electrons flow in one direction around this circuit, the positively charged metal ions flow in the opposite direction in the electrolyte.

65. 5.2.1 Half Cells & Bridge Daniell Element‘s Half Cells:
Zinc electrode (anode): Zn(s) → Zn2+(aq) + 2 e–
Copper electrode (cathode): Cu+(aq) + 2 e– → Cu(s)
simple battery: flow of electrical current from the anode to the cathode
Bridge
completion of the electric circuit: salt bridge
ionic conduction path between the anode and cathode electrolytes
porous plug that allows ion flow without electrolyte mixing
salt bridge from electrolyte saturated gel in an inverted U-tube
reduction
oxidation

66. 5.3 Electrical Potential
cell diagram (path of the electrons in the electrochemical cell)
Daniell cell:
Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)
reduced form  oxidised form
electrical potential between anode and cathode
electrochemical cell voltage
electromotive force
emf
voltmeter ca 1,1 V Zn Cu
1 m
ZnSO4
diaphragm
(pottery)
bridge containing
KCl solution
1 m
CuSO4

67. 5.4 Standard Hydrogen Electrode
Pt electrode
H2 gas
standard hydrogen electrode (SHE)
= reference potential = E0 = 0 V
H2 ñ ð 2H+ + 2e-
p = 1,01325 bar
T = 25°c
a(H+) = 1 mol / l
c(H+) = 1,235 mol / l (HCl)

68. 5.5 Metal Standard Potentials
standard hydrogen electrode
= reference potential
E0 = 0 V
metal electrode / metal salt solution
at standard conditions
= standard metal potential
M ð Mz+ + ze-
p = 1,01325 bar
T = 25°c
c(Mz+ ) = 1 mol / l
pH < 6 precipitation prevention
Pt electrode
H2 gas

69. 5.5.1 Metal Standard Potential Tables
pH-dependant

70. 5.5.2 Calvanic Corrosion Potential Chart
Galvanic Corrosion Potential Chart
K, Na, Mg, Al, Zn, Fe, Pb, Cu, Ag, Au
passivation of Al, Mg, Mn, Cr
alternative corrosion potential charts for industrial materials
cathode
least noble
corroded metals
strong oxidation affinity
negative oxidation potential
anode
most noble
protected metals
weak oxidation affinity
positive oxidation potential
Galvanic Corrosion Potential Chart
Galvanic corrosion potential is a measure of how dissimilar metals will
corrode when placed against each other in an assembly. Metals close to
one another on the chart generally do not have a strong effect on one
another, but the farther apart any two metals are separated, the stronger
the corroding effect on the one higher in the list. This list represents the
potential available to promote a corrosive reaction, however the actual
corrosion in each application is difficult to predict. Typically, the presence
of an electrolyte (eg. water) is necessary to promote galvanic corrosion.

71. 5.6 Standard Cell Potential
E°cell = E°red(cathode) – E°red(anode) = E°red(cathode) + E°ox(anode) cell diagram: Pt(s) | H2(1 atm) | H+(1 M) || Cu2+ (1 M) | Cu(s) Standard cell potential E°cell = E°red(cathode) – E°red(anode) E°cell = E°(Cu2+/Cu) – E°(H+/H2) E°cell = 0,34 V - = V = 0,34

72. 5.6.1 Free Energy in Electrochemical Cells
operation of electrochemical cells:  chemical energy  electrical energy 
Eelectr = Q·Ecell = n· F·Ecell
Wmax = Welectrical = – n·F·E°cell
free energy: maximum amount of work that can be extracted
ΔG = – n·F·Ecell
The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a chemical reaction. This energy is referred to as electrical work and is expressed by the following equation:
where work is defined as positive into the system.
A positive cell potential gives a negative change in Gibbs free energy. This is consistent with the cell production of an electric current flowing from the cathode to the anode through the external circuit. If the current is driven in the opposite direction by imposing an external potential, then work is done on the cell to drive electrolysis.[23]

73. 5.6.2 Maximum Work in Electrochemical Cells
ΔG < 0 spontaneous electrochemical reaction:  
electric current electrochemical cells batteries fuel cells
ΔG < 0 electrolysis 
ΔG = ΔG° + RT·lnK
nFΔE = nFΔE° – RT lnK
n  = the number of electrons/mole product
F  = the Faraday constant (coulombs/mole)
ΔE  = cell potential

74. 5.6.3 Exercise: Gibbs Free Energy
What happens if ΔG = 0

75. electrode potential dependency on temperature and concentration
5.7 NERNST‘s Equation 1
electrode potential dependency on temperature and concentration
E = measured cell potential
E0 = standard reaction potential
R = gas constant ( 8,3145 J . mol-1 . K-1)
T = Kelvin temperature
z = charges
F = Faraday’s constant
[ ] = concentration of oxidant / reductant in mol / l

76. ( = metal electrode in metal salt solution)
NERNST’s Equation 2
type electrode
( = metal electrode in metal salt solution)
[red] = const
E = measured cell potential
E0 = standard reaction potential
R = gas constant ( 8,3145 J . mol-1 . K-1)
T = Kelvin temperature
z = charges
F = Faraday’s constant
[ox] = concentrationen of oxidant in mol / l

77. 5.7.2 Exercise: Maximum Electrical Voltage
Calculate the maximum electrical voltage for the Daniell element when standard conditions !
Daniell Element: Cu/Cu++//Zn++/Zn
Cu/Cu++/ +0,34 V Zn++/Zn/ +0,76 V
S = + 1,1 V
Calculate the maximum electrical voltage for a galvanic cell with Ni/Ni++//Zn++/Zn when standard conditions !
Ni/Ni++// -0,23 V Zn++/Zn/ +0,76 V
S = + 0,53 V

78. 5.7.3 Exercise: Nernst Equation
What is the electrode potential for a silver electrode at 0°C when the Ag+ concentration is 1 mol ?

79. 5.8 Ag / AgCl Electrode Ag K+ Ag+ Cl- AgClsat 2. type electrode
= metal electrode in saturated metal salt solution
= electrode with constant potential (no concentration changes)
T = 25 °C:
1 m KCl E0 = + 0,220 V
sat. KCl E0 = + 0,1958 V Ag
AgClsat K+
Ag+
Cl-

80. Concentration Cells
Cu(s) | Cu2+ (0.05 M) || Cu2+ (2.0 M) | Cu(s)
half cell reactions :
oxidation: Cu(s) → Cu2+ (0.05 M) + 2 e–
reduction: Cu2+ (2.0 M) + 2 e– → Cu(s)
overall reaction: Cu2+ (2.0 M) → Cu2+ (0.05 M)
cell's emf :
E = E°- ( \2) log [0,05/2] = V
E° = 0 , (electrodes and ions are the same in both half-cells)

81. 5.10 pH Electrode measurement electrode = glas electrode
reference electrode
= silver chlorid
flask cork
KCl - solution
Ag-stick
AgClsolid
KCl-
solution
glas membran
= solid electrolyt
buffer solution

82. 5.11 Battery

83. 5.12 Electrochemical cell types

84. 5.13 Primary electrochemical cells

85. 5.14 Secondary electrochemical cells

86. 5.15 Dry Cells Leclanché's cell
moist electrolyte paste
Leclanché's cell
anode is a zinc container surrounded by a thin layer of MnO2
Cathode a carbon bar inserted on the cell's electrolyte
moist electrolyte paste NH4Cl + ZnCl2 mixed with starch
Anode: Zn(s) → Zn2+(aq) + 2 e–
Cathode: 2 NH4+(aq) + 2 MnO2(s) + 2 e– → Mn2O3(s) + 2 NH3(aq) + H2O(l)
Overall reaction:
Zn(s) + 2 NH4+(aq) + 2 MnO2(s) → Zn2+(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l)
E = ~ 1.5 V
Starch = Stärke

87. 5.16 Zn Battery
Graphics:

88. 5.17 Mercury Battery
amalgamated anode of mercury and zinc surrounded by a stronger alkaline electrolyte and a paste of ZnO and HgO Mercury battery half reactions are shown below: Anode: Zn(Hg) + 2 OH–(aq) → ZnO(s) + H2O(l) + 2 e– Cathode: HgO(s) + H2O(l) + 2 e– → Hg(l) + 2 OH–(aq) Overall reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l) no changes in the electrolyte's composition when working 1.35 V of direct current Not rechargeable
Graphics:

89. 5.18 Lead-Acid battery
six identical cells assembled in series (6 x 2V ) = 12 V lead anode lead dioxide cathode Electrolyte sulfuric acid Anode: Pb(s) + SO42–(aq) → PbSO4(s) + 2 e– Cathode: PbO2(s) + 4 H+(aq) + SO42–(aq) + 2 e– → PbSO4(s) + 2 H2O(l) Overall reaction: Pb(s) + PbO2(s) + 4 H+(aq) + 2 SO42–(aq) → 2 PbSO4(s) + 2 H2O(l) Rechargeable (external voltage  electrolysis of the products)
The lead-acid battery used in automobiles

90. 5.19 Lithium rechargeable battery (1)
Positive electrodes Electrode material Average potential difference Specific capacity Specific energy LiCoO2 3.7 V 140 mA·h/g kW·h/kg LiMn2O4 4.0 V 100 mA·h/g kW·h/kg LiNiO2 3.5 V 180 mA·h/g kW·h/kg LiFePO4 3.3 V 150 mA·h/g kW·h/kg Li2FePO4F 3.6 V 115 mA·h/g kW·h/kg LiCo1/3Ni1/3Mn1/3O2 3.6 V 160 mA·h/g kW·h/kg Li(LiaNixMnyCoz)O2 4.2 V 220 mA·h/g kW·h/kg Negative electrodes Graphite (LiC6) V 372 mA·h/g kW·h/kg Hard Carbon (LiC6) Titanate (Li4Ti5O12) 1-2 V 160 mA·h/g kW·h/kg Si (Li4.4Si)[27] V 4212 mA·h/g kW·h/kg Ge (Li4.4Ge)[28] V 1624 mA·h/g kW·h/kg

91. Lithium rechargeable battery (2)
The following equations are in units of moles, making it possible to use the coefficient x.
Overdischarge supersaturates lithium cobalt oxide, leading to the production of lithium oxide
Overcharge up to 5.2 Volts leads to the synthesis of cobalt(IV) oxide
In a lithium-ion battery the lithium ions are transported to and from the cathode or anode,
with the transition metal, cobalt (Co), in LixCoO2 being oxidized from Co3+ to Co4+
during charging, and reduced from Co4+ to Co3+ during discharge.

92. 5.20 Flow battery

93. 5.21 Fuel cells

94. 6.0 Practical Potential Charts

95. 6.1 Galvanic series (most noble at top)

96. 6.1.1 Leiterplattenherstellung

97. 7.0 Corrosion less noble metal = anode M1 ® M1z+ + z e-
noble metal = cathode M2z+ + z e- ® M2
acidic electrolyte H+ + 2 e- ® H2­
less acidic electrolyte H2O + O2 + 4 e- ® 4 OH-
different local electrolyte concentrations
corrosion from different local oxygen content

98. 7.1 Fe Corrosion in a Drop of Water
corrosion zone
Fe++
few O2
OH-
more O2
Fe OH- ® Fe(OH)2 ® Fe(OH)3

99. 7.2 Iron corrosion

100. 7.3 Corrosion of common metals

101. 7.4 Prevention of corrosion

102. 7.5 Coating

103. 7.6 Sacrificial anodes

104. 7.7 Electrolysis

105. 7.8 Electrolysis of molten sodium chloride

106. 7.9 Electrolysis of aqueous solutions

107. 7.10 Electrolysis of a solution of sodium chloride

108. 7.11 Applications

109. Exercises 1 What is the internal energy of 1 mole Ar at 0°C ?
What is the volume of 1 mole hydrogen gas at 25 °C ?
What is the entropy change in 1 mole hydrogen gas at standard conditions when increasing the volume to DV = 1 m3 ?
The equilibrium constant for acetic acid in water at 25°C is 4,76. What is Gibbs Free Energy at that temperature ?
Calculate the maximum electrical voltage for the DANIELL element when normal pressure and 10 °C !
Calculate the maximum electrical voltage for a galvanic cell with Ni/Ni++//Zn++/Zn when normal pressure and 10 °C !
Explain the difference between a galvanic and an electrolytic cell !
What is the standard hydrogen potential ?

110. Exercises 2 What is the standard metal potential ?
How can you decide whether an ion will precipitated at a given electrode ?
What is the electrode potential for a silver electrode at 10°C when the Ag+ concentration is 1 mol ?
How can you calculate the amount of elementary metal to be formed on an electrode ?
How can you calculate the maximum energy which can be obtained from a battery
Explain the chemical potential !
Explain the lead/acid battery !
Explain the mercury battery !

111. Exercises 3 Explain the pH electrode ! What is the calomel electrode?
What is the standard metal potential ?
Explain the fuel cell !
Explain corrosion processes ia a water drop !

112. Web Links http://en.wikipedia.org/wiki/Electrochemistry#Principles

113. References
A. Burrows, A. Parsons , G. Price, J. Holman , G. Pilling; Chemistry: Introducing inorganic, organic and physical chemistry ; Oxford University Press 2009
J. Hoinkins; E. Lindner; Chemie für Ingenieure; Verlag: Wiley-VCH Verlag GmbH & Co. KGaA, 2007
P.W. Attkins; L. Jobnes; Chemie – einfach alles; Verlag: Wiley-VCH Verlag GmbH & Co. KGaA, 2006
Römpp‘s Chemie Lexikon
DTV-Atlas zur Chemie