PRINSIP KOROSI - destruction of materials karena reaksi kimia dengan lingkungannya

Corrosion  destruction of materials karena
reaksi kimia dengan lingkungannya
Corrosion Science Ilmu yang mempelajari mekanisme
proses korosi, penyebabnya, dan metoda
pencegahannya
Corrosion scientist mempelajari mekanisme korosi dan
pengembangan cara2 pencegahannya,
material lain, inhibitor alternatif dll.
Corrosion engineering aplikasi ilmu korosi dan seni
untuk mencegah terjadinya korosi secara
ekonomi dan aman
Corrosion engineer

1. PRINSIP KOROSI
RINI RIASTUTI

2. INTRODUCTION
•CORROSION  DESTRUCTION OF MATERIALS KARENA
REAKSI KIMIA DENGAN LINGKUNGANNYA
•CORROSION SCIENCE ILMU YANG MEMPELAJARI MEKANISME
PROSES KOROSI, PENYEBABNYA, DAN METODA
PENCEGAHANNYA
•CORROSION SCIENTIST MEMPELAJARI MEKANISME KOROSI DAN
PENGEMBANGAN CARA
2
PENCEGAHANNYA,
MATERIAL LAIN, INHIBITOR ALTERNATIF DLL.
•CORROSION ENGINEERING APLIKASI ILMU KOROSI DAN SENI
UNTUK MENCEGAH TERJADINYA KOROSI SECARA
EKONOMI DAN AMAN
•CORROSION ENGINEER

3. DEFINITION
•CORROSION IS THE
DEGRADATION OF A MATERIAL,
USUALLY A METAL, BY CHEMICAL
OR ELECTROCHEMICAL REACTION
WITH ITS ENVIRONMENT
Or simply, Corrosion is extractive
metallurgy in reverse.

4. MATERIALS:
•METALS (IRON, ALUMINIUM, ZINC ETC)
•NONMETALLIC MATERIALS (PLASTIC, RUBBER, CERAMICS,
CONCRETE )
•DETERIORATION OF PAINT AND RUBBER BY SUNLIGHT OR
CHEMICALS

5. ENVIRONMENTS :
•PRACTICALLY ALL ENVIRONMENTS ARE CORROSIVE TO SOME DEGREE.
•EXAMPLES :
•WATERS (FRESH, DISTILLED, SALT, MINE WATERS)
•ATMOSPHERES (RURAL, URBAN, AND INDUSTRIAL ATMOSPHERES)
•GASES (CHLORINE, AMMONIA, HYDROGEN SULFIDE, SULFUR DIOXIDE, AND FUEL
GASES)
•MINERAL ACID ( HYDROCHLORIC, SULFURIC, NITRIC)
•SOILS
•VEGETABLE OIL, PETROLEUM OILS
•ETC.

6. 1.2 THE CO$T OF CORROSION
•DIRECT COST:
•REPLACEMENT OF CORRODED COMPONENTS
•USE OF CORROSION RESISTANT ALLOYS
•USE OF COATINGS OR INHIBITORS
•ELECTROCHEMICAL PROTECTION MEASURES

7. CORROSION DAMAGE
•APPEARANCE
•MAINTANANCE AND OPERATING COST
•PLANT SHUTDOWN
•CONTAMINATION OF PRODUCT
•LOST OF VALUABLE PRODUCT
•EFFECT ON SAFETY AND RELIABILITY
•PRODUCT LIABILITY

8. INDIRECT COST
INDIRECT COST RESULTING FROM ACTUAL OR POSSIBLE
CORROSION ARE DIFFICULT TO EVALUATE.
•LOSS OF PRODUCTION DURING DOWNTIME
•LOSS OF PRODUCTS DUE TO LEAKAGE
•LOSS OF EFFICIENCY DUE TO CORROSION
•CONTAMINATION
•LOSS OF HUMAN LIVES DUE TO EXPLOSION, FIRE ETC.

9. CLASIFICATION CORROSION
•WET CORROSION  LIQUID PHASE (INCL AQUEOUS
CORR , ELECTROLYTE)
•DRY CORROSION  BERHUBUNGAN DENGAN ‘HIGH
TEMPERATURE’
•WITHOUT LIQUID PHASE
•DIATAS ‘DEW POINT’ OF ENVIRONMENT

10. CORROSION PRINCIPLES
•FACTORS AFFECTING CHOICE OF ENGINEERING MATERIAL :
•MATERIAL
•AVAILABILITY
•STRENGTH
•FABRICABILITY
•APPEARANCE
•CORROSION RESISTANCE
•COST

11. •FACTORS AFFECTING CORROSION RESISTANCE OF METAL
•CORROSION RESISTANCE
•PHYSICAL CHEMICAL
•ELECTROCHEMICAL
•THERMODYNAMIC
•METALLURGICAL
* ELECTROCHEMISTRY N
THERMODYNAMIC ARE GREAT
IMPORTANCE IN UNDERSTANDING
AND CONTROLLING CORROSION

12. CORROSION RATE
•CORROSION RATE EXPRESSIONS :
- UNTUK MENGEKSPRESIKAN SECARA KUANTITATIF KECEPATAN
PENYERANGAN KOROSI UNTUK SETIAP MATERIAL.
•CR EXPRESSED IN A VARIETY OF WAYS :
•% WEIGHT LOSS
•MG PER SQ-CM PER DAY
•GRAM PER SQ-IN PER HR
•MILS PER YEAR (MPY)
•ELECTROCHEMICAL :
•POLARIZATION
•EIS

14. RUST (CORROSION PRODUCT)

15. STRESS CORROSION
“
Cross section of stress-corrosion
crack in stainless steel.
Transgranular” SCC (“TGSCC”)Transgranular” SCC (“TGSCC”)
Intergranular” SCC (“IGSCC”)Intergranular” SCC (“IGSCC”)

16. Area 1
Area 2
Area 3
Area 4
Area 5Upper platform
Stack area
Fence section
Lower platform, area
substructure

17. Corroded H beam in area 1(substructure), 8 o’clock position

21. hole

24. 1.3 CORROSION CAN BE
CONTROLLED
•MATERIALS SELECTION
•PROPER DESIGN
•ELECTROCHEMICAL PROTECTION
•INHIBITORS
•PAINTS/COATINGS

25. 1.4 SIGNIFICANCE OF
CORROSION CONTROL
•ECONOMICS
•SAFETY
•ENVIRONMENTAL CONCERNS

26. ELECTROCHEMICAL
ASPECT
PRINSIP KOROSI

27. TOPICS :
•BASIC THEORY OF AQUEOUS CORROSION
•ELECTROCHEMICAL ASPECT
•THERMODYNAMICS ASPECT
•KINETICS ASPECT

28. electrochemistry
ionic
electrodic
Electrolyte thermodynamics
Mass Transport & Conductivity
Electrode thermodynamics
Electrode Kinetics
The Division Of Electrochemistry

29. The concept of electrode
1. In corrosion
Electrodes are pieces of metal on which an
electrochemical reaction is occurring.
From definition of corrosion is the
degradation of metals resulting from
their chemical interaction with the
environment.
Example : a mild steel immersed in
acid solution
The metal in this environment is not
stable and shows a natural tendency
toward spontaneous oxidation.

30. ELECTROCHEMICAL NATURE OF
AQUEOUS CORROSION
•METALLIC CORROSION PROCESS INVOLVE TRANSFER OF ELECTRONIC
CHARGE IN AQUEOUS SOLUTIONS
•THUS, IT IS NECESSARY TO DISCUSS ‘THE ELECTROCHEMICAL NATURE
OF CORROSION’ BEFORE DISCUSSING OF THE VARIOUS FORM OF
CORROSION.

31. ELECTROCHEMICAL REACTIONS :
EXAMPLE OF CORROSION BETWEEN FE AND HYDROCHLORIC ACID,
REPRESENTED BY :
FE + 2 HCL  FECL
2
+ H
2
(1)
FE REACTS WITH THE ACID SOLUTION FORMING SOLUBLE FE CHLORIDE
AND LIBERATING HYDROGEN BUBBLES ON THE SURFACE.

32. ACIDS AND BASES
•AN ACID IS A SUBSTANCE THAT PRODUCES EXCESS HYDROGEN IONS
(H
+
) WHEN DISSOLVED IN WATER
•EXAMPLES ARE HCL, H
2
SO
4
•A BASE IS A SUBSTANCE THAT PRODUCES EXCESS HYDROXYL IONS
(OH
-
) WHEN DISSOLVED IN WATER
•EXAMPLES ARE NAOH, KOH

33. CORROSION OF ZINC IN ACID
•ZINC DISSOLVES WITH HYDROGEN EVOLUTION
ZN + 2HCL  ZNCL
2 + H
2

Zinc known as a base or active metalZinc known as a base or active metal
One atom of zinc metalplus two molecules of hydrogen
chloride (hydrochloric acid)
reacts to form
goes to
one molecule of zinc chlorideplus one molecule of hydrogen gas

35. ANODIC REACTIONS
•EXAMPLES
ZN ZN
2+
+ 2E
-
ZINC CORROSION
FE FE
2+
+ 2E
-
IRON CORROSION
AL AL
3+
+ 3E
-
ALUMINIUM CORROSION
FE
2+
 FE
3+
+ E
-
FERROUS ION OXIDATION
H
2 2H
+
+ 2E
-
HYDROGEN OXIDATION
2H
2
O O
2
+ 4H
+
+ 4E
-
OXYGEN EVOLUTION
•OXIDATION REACTIONS
•PRODUCE ELECTRONS

36. CATHODIC REACTIONS :
HYDROGEN EVOLUTION 2H
+
+ 2E  H
2
OXYGEN REDUCTION (ACID) O
2 + 4H
+
+ 4E  2H
2O
OXYGEN REDUCTION(BASIC): O
2+ 2H
2O+4E4OH
-
METAL REDUCTION : M
3+
+ E  M
2+
METAL DEPOSITION ; M
+
+ E  M
read Fontana p 17

37. ELECTROCHEMISTRY OF CORROSION
•PERMUKAAN LOGAM YG
SAMA DIEKSPOSE DI
AQUEOUS ELECTROLYTE
•BIASANYA MEMILIKI SISI
UNTUK SUATU OKSIDASI
DAN REDUKSI. SISI INI
MEMBENTUK CORROSION
CELL

38. THE ELECTRODE - SOLUTION INTERFACE IS ALSO HYDRATED BECAUSE OF THE
CHARGE EFFECTS.

( N.B. THE POTENTIAL DIFFERENCE BETWEEN METAL AND SOLUTION IS OF THE
ORDER OF  1V, REMEMBER, - CANNOT BE MEASURED - BUT OVER A
MOLECULAR DISTANCE AMOUNTS TO AN ELECTRIC FIELD OF  10
8
V/CM).
Schematic of a charged
interface and the locations
of cations at the electrode
surface.

39. BECAUSE OF ADSORBED H
2
O ON SURFACE AND PRIMARY HYDRATION SHELL, IONS
CAN ONLY APPROACH TO WITHIN A FIXED DISTANCE FROM THE ELECTRODE:
THE OUTER HELMHOLTZ PLANE. THE RESULTING DOUBLE LAYER OF CHARGE IS
OFTEN TREATED THEORETICALLY AS A CAPACITOR . . . .
Simplified double layer at a metal aqueous interface.

40. VOLTAIC CELLS
IN SPONTANEOUS
OXIDATION-REDUCTION
(REDOX) REACTIONS,
ELECTRONS ARE
TRANSFERRED AND
ENERGY IS RELEASED.

41. VOLTAIC CELLS
•IN THE CELL, THEN,
ELECTRONS LEAVE THE
ANODE AND FLOW
THROUGH THE WIRE
TO THE CATHODE.
•AS THE ELECTRONS
LEAVE THE ANODE, THE
CATIONS FORMED
DISSOLVE INTO THE
SOLUTION IN THE
ANODE
COMPARTMENT.

42. VOLTAIC CELLS
•AS THE ELECTRONS
REACH THE CATHODE,
CATIONS IN THE
CATHODE ARE
ATTRACTED TO THE
NOW NEGATIVE
CATHODE.
•THE ELECTRONS ARE
TAKEN BY THE
CATION, AND THE
NEUTRAL METAL IS
DEPOSITED ON THE
CATHODE.

43. ELECTROMOTIVE FORCE (EMF)
•WATER ONLY
SPONTANEOUSLY FLOWS
ONE WAY IN A
WATERFALL.
•LIKEWISE, ELECTRONS
ONLY SPONTANEOUSLY
FLOW ONE WAY IN A
REDOX REACTION—FROM
HIGHER TO LOWER
POTENTIAL ENERGY.

44. ELECTROMOTIVE FORCE (EMF)
•THE POTENTIAL DIFFERENCE BETWEEN THE ANODE AND CATHODE IN
A CELL IS CALLED THE ELECTROMOTIVE FORCE (EMF).
•IT IS ALSO CALLED THE CELL POTENTIAL, AND IS DESIGNATED E
CELL.

45. CELL POTENTIAL
CELL POTENTIAL IS MEASURED IN VOLTS (V).
1 V = 1
J
C

46. STANDARD REDUCTION POTENTIALS
REDUCTION
POTENTIALS
FOR MANY
ELECTRODES
HAVE BEEN
MEASURED AND
TABULATED.

47. STANDARD HYDROGEN ELECTRODE
•THEIR VALUES ARE REFERENCED TO A STANDARD
HYDROGEN ELECTRODE (SHE).
•BY DEFINITION, THE REDUCTION POTENTIAL FOR
HYDROGEN IS 0 V:
2 H
+
(AQ, 1M) + 2 E
−
 H
2 (G, 1 ATM)

48. STANDARD CELL POTENTIALS
THE CELL POTENTIAL AT STANDARD CONDITIONS CAN BE FOUND
THROUGH THIS EQUATION:
E
cell= E
red
(cathode) − E
red
(anode) 
Because cell potential is based on the
potential energy per unit of charge, it
is an intensive property.

49. CELL POTENTIALS
•FOR THE OXIDATION IN THIS CELL,
•FOR THE REDUCTION,
E
red
= −0.76 V
E
red
= +0.34 V

50. CELL POTENTIALS
E
cell=E
red(cathode) − E
red(anode)
= +0.34 V − (−0.76 V)
= +1.10 V

51. OXIDIZING AND REDUCING AGENTS
•THE STRONGEST OXIDIZERS
HAVE THE MOST POSITIVE
REDUCTION POTENTIALS.
•THE STRONGEST REDUCERS
HAVE THE MOST NEGATIVE
REDUCTION POTENTIALS.

52. OXIDIZING AND REDUCING AGENTS
THE GREATER THE DIFFERENCE
BETWEEN THE TWO, THE GREATER
THE VOLTAGE OF THE CELL.

53. FREE ENERGY
G FOR A REDOX REACTION CAN BE FOUND BY USING THE
EQUATION
G = −NFE
WHERE N IS THE NUMBER OF MOLES OF ELECTRONS TRANSFERRED,
AND F IS A CONSTANT, THE FARADAY.
1 F = 96,485 C/MOL = 96,485 J/V-MOL

54. CORROSION AND…

55. …CORROSION PREVENTION

56. FARADAY’S LAW
•CHARGE IS RELATED TO MASS OF MATERIAL REACTED IN AND
ELECTROCHEMICAL REACTION:
2H
+
+ 2E
-
 H
2
Two hydrogen
ions
React with two
electrons
To produce one molecule of
hydrogen gas

57. FARADAY’S CONSTANT
•ONE MOLE OF HYDROGEN IONS (1 G) CONTAINS AVOGADRO’S
NUMBER (6 10
23
) IONS
•HENCE ELECTRONS WILL REACT WITH EACH MOLE OF HYDROGEN
IONS
•CHARGE ON THE ELECTRON IS 1.6  10
-19
C
•HENCE ONE MOLE OF IONS REQUIRES 96500 C
•THIS IS KNOWN AS FARADAY’S CONSTANT

58. FARADAY’S LAW
(g/mole) metal of weight atomic
(g) oxidised metal of mass
ed transferrelectrons ofnumber
C/mole) (96500constant sFaraday'
(C) charge where






M
m
n
F
Q
M
nFm
Q

59. EFFECT OF POTENTIAL
•ELECTROCHEMICAL REACTIONS INVOLVE TRANSFER OF CHARGE
•HENCE, WE EXPECT THAT THE VOLTAGE OF THE METAL WITH RESPECT
TO THE SOLUTION WILL AFFECT ELECTROCHEMICAL REACTIONS
•VOLTAGE OF METAL WITH RESPECT TO SOLUTION IS KNOWN AS THE
ELECTROCHEMICAL POTENTIAL

60. AQUEOUS CORROSION THERMODYNAMICS
THE CORROSION OF METALS IN CONTACT WITH AQUEOUS SOLUTIONS INVOLVES
ELECTRONIC AND IONIC PROCESSES AT SURFACES . . . I.E., IT INVOLVES . . .
(1) ELECTRODE PROCESSES
ELECTRONATION . . . M(AQ)
2+
+ 2E  M(S)
(CATHODIC)
DE-ELECTRONATION . . .M(S)  M(AQ)
2+
+ 2E
(ANODIC)
THESE ARE COUPLED PROCESSES . . . THE OXIDIZED SPECIES PROVIDE THE ELECTRONS
FOR THE REDUCED SPECIES.
(2) THE REACTIONS OCCUR AT THE METAL-SOLUTION INTERFACE
60
UN1001: Section 12: Aqueous Corrosion ThermodynamicsUN1001: Section 12: Aqueous Corrosion Thermodynamics

61. PREDICTION OF
CORROSION
TENDENCIES
THERMODYNAMIC ASPECT

62. TERMODINAMIKA KOROSI
•KOROSI : TENDENSI LOGAM  OKSIDANYA
•REAKSI OKSIDASI  MEMBEBASKAN ENERGI
•TENDENSI OKSIDASI LOGAM TERGANTUNG POTENSIAL LOGAM
•E
EQ
= E
REDUKSI
- E
OKSIDASI

63. KESTABILAN TERMODINAMIK CU DAN CU
++
Cu stabil dalam
bentuk logam
Cu stabil dalam bentuk ion
E
cu++/Cu
active
nobel

64. AKAN DIBAHAS :
1.ASPEK TERMODINAMIKA UNTUK MEMAHAMI ELEKTROKIMIA KOROSI
2.PERSAMAAN NERNST
UNTUK MENGHITUNG POTENSIAL
KESETIMBANGAN  UNTUK MEMPREDIKSI
DAERAH KOROSI DAN KESTABILAN -
DIPERLUAS BAHASAN UNTUK DIAGRAM EH -
PH

65. ELEKTRODA REVERSIBLE DAN
PERSAMAAN NERNST
•ELEKTRODA REVERSIBLE
•CIRI : REAKSI ELEKTROKIMIA DAPAT DI BOLAK
BALIK DENGAN MEMBALIKAN ARAH ARUS LISTRIK
CONTOH ;
•ELEKTRODA CU++/CU.
•JIKA BERTINDAK SEBAGAI ANODA, CU  CU++ + 2E. DENGAN
MEMBALIKAN ARUS, INI MENJADI KATODA : CU++ + 2E  CU

66. PELAT BESI DIDALAM LARUTAN
ASAM
PROSES ANODIC :
FE  FE++ + 2E ,
DAN BILA ARUS DIBALIK ARAH MAKA YANG AKAN TERJADI
2H+ + 2E  H2.

67. 2.CIRI LAIN :
ELEKTRODA REVERSIBLE YANG IDEAL
DICIRIKAN OLEH ‘FREE ENERGY” ,
ΔG = ENERGI KEADAAN AKHIR -ENERGI
AWAL,
ANTARA PRODUK DAN REAKTAN DARI SUATU
REAKSI ELEKTROKIMIA.
DENGAN KATA LAIN,
GOKSIDASI = - G REDUKSI

68. •UNTUK REAKSI ELEKTROKIMIA UMUM :
OX + NE  RED
•ENERGY BEBAS DIHUBUNGKAN DENGAN REDUKSI
REVERSIBLE DIBERIKAN OLEH PERSAMAAN :
DGRED = GRED - GOKS ,
PADA KASUS OKSIDASI MENGIKUTI
DGOKS = GOKS – GRED.

69. •PADA SYSTEM ELEKTROKIMIA PADA T,P KONSTAN, ENERGY YANG
TERLIBAT ADALAH PERUBAHAN FREE ENERGY , ΔG.
•ΔG INI DAPAT DIHUBUNGKAN DG POTENSIAL ELEKTRODA
REVERSIBLE DENGAN PERSAMAAN BERIKUT :
ΔG = - NFE (II-1)
INI MERUPAKAN DASAR PERHITUNGAN POTENSIAL.

70. •REAKSI ELEKTROKIMIA UMUM, DITULIS DALAM ARAH KONVENSIONAL :

OX + NE  RED.
•ENERGI BEBAS YANG TERLIBAT DALAM PROSES :
DG= G
RED
(PRODUK) - G
OKS
(FREE ENERGY
REACTANT)

71. PERSAMAAN TERMODINAMIKA
UMUM UNTUK KASUS INI DAPAT
DITULIS
)2()(ln
)(ln


II
oks
red
RTG
oks
red
RTGGG
o
o
oks
o
red

72. •PERS II-1 DIPAKAI UNTUK MENGEKSPRESIKAN POTENSIAL:
nF
G
E



73. •PERS II-1 DAN II-2 MEMBERIKAN NERNST EQ
)3(][ln  II
red
ox
nF
RT
EE
o
R= konstanta gas
T= temperature (Kelvin)
N= jumlah electron yang terlibat dalam
reaksi elektrokimia
F= konstanta Faraday
(ox)= aktivitas zat teroksidasi ( gunakan
konsentrasi zat teroksidasi)
(red) = aktivitas zat tereduksi

E
o
= potensial standar suatu elektroda.

74. NERNST EQUATION
DIPAKAI UNTUK :
•MENENTUKAN CORROSION TENDENCY DARI
BEDA LOGAM.
•MENGHITUNG BEDA POTENSIAL PADA KOROSI
BIMETALLIC DAN
• MENGHITUNG ENERGY MAKSIMUM YANG ADA
DARI SEL GALVANIC.
•DIAGRAM E – PH JUGA DAPAT DIHITUNG YANG
MENUNJUKKAN KONDISI UNTUK KESTABILAN
LOGAM DALAM MEDIA AQUEOUS DAN KONDISI
APAKAH LOGAM TIDAK STABIL DANMEMILIKI
KECENDERUNGAN UNTUK KOROSI.

75. •CONTOH, ELEKTRODA COPPER CU++/CU,
DENGAN REAKSI ELEKTROKIMIA
CU++ + 2E  CU.
•AKTIVITAS ZAT MURNI = 1. MAKA, [CU] = 1. PERSAMAAN NERNST
DAPAT DITULIS DENGAN CARA :
)4(][ln/
/


 IICu
nF
RT
EE CuCu
o
CuCu

76. •NILAI POTENSIAL REVERSIBLE TERGANTUNG PADA
KONSENTRASI ION2 DALAM LARUTAN, DAN ECU++/CU
MENINGKAT DENGAN KONSENTRASI CU++ DALAM ELEKTROLIT
MENINGKAT.

77. •. PADA TABLE INI DAPAT DILIHAT BAHWA
E
O
CU++/CU = + 0.34 V (II-5)
PENTING DISINI UNTUK MENEKANKAN KONVENSI REDUKSI
SEBAGAI DASAR PERSAMAAN NERNST SPT PERSAMAAN II-3,

78. •PADA PERSAMAAN II-5, PERSAMAAN NERNST UNTUK COPPER
YANG DICELUP DALAM LARUTAN YANG MENGANDUNG CU++
DAPAT DITULIS DENGAN CARA BERIKUT, MENGGUNAKAN LOG :
)(log
3.2
34.0
Cu
Cu
nF
RT
E



79. •KOEFISIEN LOG MUDAH DIHITUNG PADA 25
O
C
•DENGAN, R = 1.987 CAL/DEG. MOLE
F = 23,060 CAL/V EQU
2.3 RT/F = 0.06
•INI MEMBERIKAN PERSAMAAN NERNST PADA 25
O
C
)(log
2
06.0
34.0
/
Cu
Cu
E
Cucu

 

80. •PADA PERSAMAAN (II-6) , NILAI N = 2, KARENA
REAKSI CU++ + 2E  CU, MELIBATKAN 2 ELEKTRON.
NILAI [CU] ADALAH AKTIVITAS COPPER MURNI, = 1.
•NILAI POTENSIAL ECU
++
/CU TETAP TERGANTUNG
PADA KONSENTRASI [CU
++
] DALAM LARUTAN, DAN
UNTUK [CU
++
] = 10
-2
G ION/L
ECU++/CU = 0.28 V
•INI HANYA JIKA [CU
++
] = 1 = 10
O
MAKA
E CU
++
/CU = E
O
CU++/CU = + 0.34V

81. ARTI POTENSIAL REVERSIBEL
•NERNST  MENGHITUNG POTENSIAL REVERSIBLE
INFORMASI TENTANG KOROSI.
•CONTOH, JIKA PELAT COPPER DI CELUP DALAM LARUTAN
YANG MENGANDUNG ION CU++, POTENSIAL REVERSIBLE
MEMBERIKAN NILAI POTENSIAL ELEKTRODA YANG
BERHUBUNGAN DENGAN SUATU KESETIMBANGAN ANTARA
CU++ DAN LOGAM CU.
COPPER, TERIONISASI DAN TERREDUKSI,  AKAN STABIL
PADA KONDISI INI SECARA BERSAMAAN.
•UNTUK KONSENTRASI 0.01 G ION/L  POTENSIAL REVERSIBLE
NYA = + 0.28 V.
•NILAI INI AKAN DITUNJUKKAN DALAM SKALA POTENSIAL
(TABEL).

82. •KOROSI LOGAM TERJADI JIKA POTENSIALNYA MEMILIKI NILAI
LEBIH MULIA DARIPADA POTENSIAL REVERSIBELNYA.
•JIKA BERLAWANAN, POTENSIAL COPPER, MISALNYA, DIBAWA KE
SUATU NILAI LEBIH AKTIF DARIPADA POTENSIAL
REVERSIBELNYA, PERSAMAAN NERNST MENUNJUKKAN
KELEBIHAN CU++.
MAKA KEMUDIAN AKAN TERJADI REAKSI REDUKSI.
REAKSI NYA CU++ + 2E  CU.

83. PENGARUH KONSENTRASI ION PADA
KECENDERUNGAN KOROSI
•PADA KASUS ELEKTRODA CU++/CU PADA
25
O
C , PERSAMAAN NERNST DAPAT DITULIS
•INI MEMUNGKINKAN UNTUK MENGHITUNG POTENSIAL
REV UNTUK BERBAGAI KONSENTRASI ION DALAM
LARUTAN.
•KONSENTRASI SERING DIPAKAI DARI TABEL YANG
DIBERIKAN TABEL II-1 UNTUK ELEKTRODA CU+=/CU.
)(log
2
06.0
34.0
/

  CuE
Cucu

84. TABLE II-1
COPPER ELECTRODE REVERSIBLE POTENTIAL AS
FUNCTION OF CU++ CONCENTRATION
_______________________________________________
CU++CONCENTRATION ECU++/CU CONCENTRATION
(G ION/L) NOTATION
____________________________________________
1 0.34 0
10-2 0.28 2
10-4 0.22 4
10-6 0.16 6
_______________________________________________

85. LATIHAN
•A ZINC PLATE IS INTRODUCE INTO A PH=1 SOLUTION.
CAN YOU PREDICT WHETHER OR NOT THE METAL WILL CORRODE ?
EXPLAIN YOUR ANSWER AND WRITE THE ELECTRODE REACTIONS

86. Potential-pH Diagrams (“Pourbaix Diagrams”)
Graphical representations of the domains of stability of
metal ions, oxides, hydroxides, etc. in aqueous solution.
potential (E)
pH

87. The Nernst Equation allows us to compute lines on the diagram for
equilibrium reactions of interest . . .
(1) Electrochemical reactions of pure charge transfer (horizontal lines -
since no H
+
or OH
-
dependence - corresponding to potentials of equilibria at
given concentrations);
(2) Pure acid-base reactions (vertical lines - since no electron transfer and
no dependence on potential - corresponding to equilibrium concentrations
of H
+
(OH
-
) for given concentrations of species);
(3) Electrochemical reactions involving charge transfer and H
+
(OH
-
)
(sloping lines).

88. (1) Pure Charge Transfer Reactions
Consider : Ni
2+
+ 2e
-
↔ Ni
so the Nernst Equation becomes:

[since, by definition, activity solid Ni = 1.0]
with E = - 0.25 V so... E = - 0.25 + 0.03log(Ni
2+
)
)ln(
2
2
 Ni
F
RT
EE
o
We usually compute for 4 conc
n
. . . 10
0
(i.e., 1.0), 10
-2
, 10
-4
, 10
-6
g-ion/L.
Partial E-pH diagram for the
Ni
2+
+ 2e
-
 Ni reaction.

89. 89
(2) Pure Acid-Base Reactions
Consider :
There is no charge transfer, so the E-pH diagram is a vertical line for a given Ni
2+

concentration.
We evaluate the equilibrium constant, K, from the standard free energy change for
the reaction :
or

KRTG
o
ln
RT
G
K
o
303.2
)log(



 HOHNiOHNi 2)(2
22
2

90. calculated from the standard chemical potentials (μ
o
) of the reactants and
products....
so
μ
o
values are tabulated . . . = -11,100 cal/mol
= -56,690 cal/mol
= -108,300 cal/mol
and = 0 (convention),

so, at 25
o
C log K = -12
o
G
 
o
RR
o
PP
o
G 
RT
K
o
H
o
OHNi
o
OH
o
Ni
303.2
2
log
22
2
)(
 


o
Ni
2
o
OH
2

o
OHNi
2
)(
o
H


91. Remember reaction

Since the continuous phases H
2
O and (solid) Ni(OH)
2
have activities of 1

or

 HOHNiOHNi 2)(2
22
2
2
2
2
2
2
]][[
]][)([
OHNi
HOHNi
K



12
][
][
log
2
2



Ni
H
]log[12]log[
22 
 NiH
]log[5.06
2
 NipH

92. AGAIN, WE COMPUTE FOR [NI
2+
] = 10
0
, 10
-2
, 10
-4
AND 10
-6
M.
92
The region to the left of a particular line (i.e., more acid) is the region of
stability for Ni
2+
. . . so, if we are on the equilibrium line at pH 9 for a Ni
2+

concentration of 10
-6
mol., and we drop the pH to 7 (say), the Ni(OH)2 solid
in contact with the solution will dissolve to try and make the concentration
to 10
-2
mol.

93. (3) REACTIONS INVOLVING ELECTRONS AND H+
CONSIDER . . .
SO THE NERNST EQUATION IS . . .



[SINCE, BY DEFINITION, [NIO] = [NI] = [H
2O] = 1]
CALCULATE E
O



[SINCE, BY DEFINITION,  ELEMENT = 0]
V


E = 0.11 -0.06 PH
93
]][[
]][[
ln
2
2
2
OHNi
HNiO
F
RT
EE
o


2
]log[03.0

 HEE
o
FnF
G
E
o
OH
o
Ni
o
NiO
o
o
2
2
 



11.0
2
56690051610



F
E
o

 eHNiOOHNi 22
2

94. PARTIAL E-PH DIAGRAM FOR THE REACTION.

 eHNiOOHNi 22
2
94
Above the line, solid NiO is stable.
Below the line, Ni metal is stable.

95. 95
NOTE : We can also depict the Ni - H2O reaction as . . .
so that
(as for NiO).
with V

E = 0.11 – 0.06pH
This is identical with the line for the Ni  NiO equilibrium . . .
i.e., Ni(OH)2 is as likely (thermodynamically) as NiO.

 eHOHNiOHNi 22)(2
22
2
2
2
2
]][[
]][)([
ln
2 OHNi
HOHNi
F
RT
EE
o


2
]log[03.0

 HEE
o
11.0
2
2
22
)(





FnF
G
E
o
OH
o
Ni
o
OHNi
o
o


96. ]][[
])([
2
2
OHNiO
OHNi
K
96
Alternatively.…
NiO + H2O = Ni(OH)2
for which
So that ∆G
o
= 0 and log K = 0 then K = 1
Since and [H2O] = 1
then [Ni(OH)2] = [NiO]
]][[
])([
2
2
OHNiO
OHNi
K
RT
o
OHNi
o
NiO
o
OH
303.2
22
)( 

RT
K
PPRR
303.2
log
 

KRTG
o
ln
RT303.2
1083005161059690 


97. NOTE THE STABILITY LINES FOR H
2O
(A)
(B)
97
E-pH diagram for nickel.
222 HeH 


 eHOOH 22
2
1
22

98. CONSIDER STABILITY OF WATER …
H
2O H
+
+ OH
-
AND [H
+
][OH
-
] = K
NOW,
AND
SO
98
RT
G
K
o
303.2
log


00
RRPP
o
G  
 
RT
K
OHOHH
303.2
log
000
2
 



99. SINCE
AT 25
O
C
I.E. [H
+
][OH
-
] = 10
-14
IN NEUTRAL SOLUTION; [H
+
] = [OH
-
]
SO = 10
-7
= -7
= 7 = PH
99
molcal
OH
/690,56
0
2

 
01.14
1363
5669037595
log 

K
0
0


H

molcal
OH
/595,37
0


Cneutral
oH
25,
][

Cn
oH
25,
]log[

Cn
oH
25,
][
1
log


100. THE REDUCTION/OXIDATION OF WATER ..
2H
+
+ 2E
-
 H
2
(2H
2O + 2E
-
 H
2 + 2OH
-
)
2H
2O  O
2 + 4H
+
+ 4E
-
FOR WHICH
AND
SINCE
AND ∆
O
= ∑ΝΜ
O

≡
∆ G
O

100
}
reduction
oxidation
K
nF
RT
E
o
ln
K
nF
RT
log
303.2

nF
o


2
2
4
2
][
]][[
log
4
303.2
OH
HO
F
RT
EE
o
OxOx


][
][
log
2
303.2
2
2
ReRe
H
H
F
RT
EE
o
dd




101. 101
 
0
2
)0)0*2((
2
2
2
Re 





FF
E
o
H
o
Ho
d

 
F
E
o
H
o
O
o
OHo
Ox
4
42
22



 
F4
)0*4(0))690,56(*2( 

V
coulomb
sampvolt
coulomb
swatt
F
coulomb
joule
F
molcoulomb
molcal
F
23.1
500,96*4
196,476
4
196,476
4
2.4*380,113
/
/
4
380,113








102. 102
]log[
4
303.2
]log[
303.2
23.1
2
O
F
RT
H
F
RT
V 

2
2
4
2
][
]][[
log
4
303.2
23.1
OH
HO
F
RT
VE
ox



 eHOOHfor 442
22
2
log0147.00591.023.1
Oox
ppHVE 
2
22 HeHfor 

][
][
log
2
303.2
0
2
2
Re
H
H
F
RT
E
d


]log[
2
303.2
]log[
303.2
2
H
F
RT
H
F
RT


2
log0295.00591.0
Re Hd
ppHE 

103. REDOX NEUTRALITY … WHEN
FROM ABOVE EQUATIONS, AT PH 7, REDOX NEUTRALITY IS AT 0.40V …
103
22
2
OH
pp
Below a, for 1 atm H2,
water decomposes to
H2
Above b, for 1 atm
O2, water
decomposes to O2

104. EQUILIBRIA FOR PARTIAL PRESSURES < 1 ATM.
104
N.B.
ERed = - 0.0591pH + 0.0295rH
EOx = 1.23 - 0.0591pH -
0.0147rO
2
2
2
2
log
1
log
log
1
log
O
O
H
H
p
p
rO
p
p
rH



105. THE NERNST EQUATION GIVES . . .
BY DEFINITION...
WE USUALLY CONSIDER CASE WHEN = 1 ATM
LINE (A) ON POURBAIX DIAGRAM . . .
105
2
22
2
//
][
log
2
059.0
H
o
HHHH
p
H
EE

 
0
2/

o
HH
E
22
log
2
059.0
059.0
/ HHH
ppHE 
2Hp
pHE
HH
059.0
2/



106. THE POURBAIX (E-PH) DIAGRAM
P
o
t
e
n
t
i
a
l
H
2
O is stable
H
2
is stable
7 14
pH = - log [H
+
]
2H
+
+ 2e
-
= H
2

Equilibrium
potential falls as
pH increases
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
2H
2O = O
2 + 4H
+
+ 4e
-
Equilibrium potential
falls as pH increases
O
2 is stable

107. IF WE HAVE AN ELECTRODE (SAY PT) IN WATER AT PH 7 EQUILIBRATED WITH H
2 AT 1
ATM., THE POTENTIAL OF THE PT WILL BE  - 400 MV. IF WE THEN LOWER THE
POTENTIAL (TO -600 MV, SAY) WATER WILL DECOMPOSE AND LIBERATE H
2 TO TRY
AND ACHIEVE A NEW EQUILIBRIUM AT A HIGHER H
2 PARTIAL PRESSURE AND/OR...
HIGHER PH.
107
The water E-pH diagram at 1 atm.

108. SIMILARLY . . . ON H
2O STABILITY LINE AT PH 7 AND  +800 MV (IN
EQUILIBRIUM WITH 1 ATM. O
2
) AND WE RAISE THE POTENTIAL, WATER WILL
DECOMPOSE AND LIBERATE O
2 TO TRY AND ACHIEVE NEW EQUILIBRIUM AT A
HIGHER O
2 PARTIAL PRESSURE AND/OR LOWER PH.
108
The water E-pH diagram at 1 atm.

109. USE OF POURBAIX DIAGRAMS
CONSIDER NI METAL CORRODING IN AN ACID SOLUTION (PH 0) CONTAINING 10
-4
M
(STRICTLY... G.ION/L) NI
2+
IN EQUILIBRIUM WITH 1 ATM. H
2
GAS.
109
Potential of the anodic regions = -0.37V
Potential of the cathodic regions = 0 V
Since the Ni  Ni
2+
reaction is more
active than the H
+
 H2
reaction, the former can drive the latter
. . . i.e., Ni will displace H2 from
solution. The surface will attain a
mixed potential EM between the two
equilibrium potentials, and corrosion
will proceed.
If, now, pH raised to  6-8 (say pH7)
then Ni  Ni
2+
is more noble than H2
evolution reaction, will not corrode via
H
+
reduction.
E-pH diagram for nickel.

110. NOTE: IF SOLUTION ALSO CONTAINS DISSOLVED O
2
AT 1 ATM, THEN AN ADDITIONAL
CATHODIC REACTION IS:
O
2
+ 2H
+
+ 2E
-
 H
2
O
ALONG LINE (B): THIS IS ALSO POSSIBLE AT PH 6-8, BUT WILL BECOME SLOWER AT HIGHER
PH, SINCE NOT MUCH H
+
IS AVAILABLE. A MIXED POTENTIAL WILL BE ESTABLISHED.
AT PH 8-14, STABLE SPECIES ARE SOLID NI(OH)
2, NI
3O
4 . . . CORROSION WILL TEND TO
PRODUCE PROTECTIVE LAYERS. . . METAL PASSIVATES.
AT PH > 14 , CORROSION (STABLE SPECIES IONIC HNIO
2
-
AT QUITE LOW POTENTIALS).
WHAT HAPPENS IF WE PUT NI METAL INTO ACID SOLUTION (SAY PH 0 IN EQUILIBRIUM WITH
1 ATM H
2) CONTAINING ZERO CONCENTRATION NI
2+
?
DISCUSS
REMEMBERING NERNST EQUATION . . .
ALTHOUGH POTENTIAL VERY LOW INITIALLY (E  - ), SOME NI
2+
GENERATED IN LIQUID
FILM NEXT TO METAL . . . POTENTIAL RISES AS [NI
2+
] INCREASES.
110

111. NOTE:
FOR CORROSION CONSIDERATIONS, A METAL ION CONCENTRATION OF 10
-6
M (10
-
6
G-ION/L) IS CONSIDERED AS INDICATIVE OF CONDITIONS THAT APPLY IN THE LIQUID
FILM NEXT TO THE METAL SURFACE WITH NO EXTRANEOUS SOURCE OF THOSE IONS . . .
MOST POURBAIX DIAGRAMS CONSIDER THIS CONCENTRATION FOR CORROSION.
POURBAIX DIAGRAMS CAN BE COMPLEX . . .
BECAUSE THERE ARE MANY
REACTIONS THAT CAN BE
CONSIDERED (NOT ALL OF THEM
OCCURRING)...
111

112. BUT, FOR INDICATING CORROSION TRENDS, WE CAN USE SIMPLIFIED VERSIONS . . .
112
A simplified form of the Pourbaix
Diagram for iron.

113. AND FOR A REACTIVE - BUT USUALLY PASSIVATED - METAL SUCH AS AL . . .
113

114. POURBAIX DIAGRAM FOR IRON
P
o
t
e
n
t
i
a
l
7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
Fe metal stable
Fe
3+
Fe oxides
stable
Will iron
corrode in
acid?
Fe
2+
stable
Yes - there is a
reasonably wide
range of potentials
where hydrogen
can be evolved and
iron dissolved
Will iron
corrode in
neutral waters?
Yes - although iron can
form an oxide in neutral
solution, it tends not to
form directly on the
metal, as the potential
is too low, therefore it is
not protective.
Will iron corrode
in alkaline
solution?
No - iron forms a solid
oxide at all potentials,
and will passivate

115. P
o
t
e
n
t
i
a
l
7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
POURBAIX DIAGRAM FOR ZINC
Equilibrium for
Zn  Zn
2+
+ 2e
-
Zn metal stable
Zn
2+
stable
in solution
Equilibrium for
Zn
2+
+ 2OH
-
 Zn(OH)
2
Zn(OH)
2
stable
solidEquilibrium for
Zn + 2OH
-
 Zn(OH)
2
+ 2e
-
Equilibrium for
Zn(OH)
2
+ 2OH
-
 ZnO
2
2-
+ 2H
2
O
ZnO
2
2-
stable in
solution
Equilibrium for
Zn + 4OH
-
 ZnO
2
2-
+ 2H
2
O + 2e
-

116. P
o
t
e
n
t
i
a
l
7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
POURBAIX DIAGRAM FOR ZINC
Zn metal stable
Zn
2+
stable
in solution
Zn(OH)
2
stable
solid
ZnO
2
2-
stable in
solution
Corrosion
C
o
r
r
o
s
io
n
Immunity
P
a
s
s
iv
it
y
Corrosion
possible with
oxygen
reduction
Corrosion
possible with
hydrogen
evolution
Corrosion requires
strong oxidising
agent
Corrosion is
thermodynamically
impossible
Corrosion is
possible, but likely
to be stifled by solid
corrosion product

117. POURBAIX DIAGRAM FOR COPPER
P
o
t
e
n
t
i
a
l
7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
Cu metal stable
Cu
2+
stable
in solution
Cu oxides
stable
C
u
O
2
2
-
s
t
a
b
le

in

s
o
ln
.
Will copper
corrode in
acid?
No - hydrogen
evolution only
occurs below the
potential for copper
corrosion
Will copper
corrode in
neutral waters?
Usually it will just
passivate, but
corrosion can occur
in slightly acid
solutions

118. POURBAIX DIAGRAM FOR ALUMINIUM
P
o
t
e
n
t
i
a
l
7 14
1.2
0.8
0.0
0.4
-1.2
-0.4
-0.8
-2.4
-1.6
-2.0
0
Al
Al
3+
Al
2O
3
AlO
2
-

119. POURBAIX DIAGRAM FOR GOLD
P
o
t
e
n
t
i
a
l
7 14
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
0
Gold metal stable
Immunity
C
C
Passivity
Gold can’t corrode
with oxygen reduction
or hydrogen evolution

120. LIMITATIONS OF POURBAIX DIAGRAMS
•TELL US WHAT CAN HAPPEN, NOT NECESSARILY WHAT WILL HAPPEN
•NO INFORMATION ON RATE OF REACTION
•CAN ONLY BE PLOTTED FOR PURE METALS AND SIMPLE SOLUTIONS,
NOT FOR ALLOYS