3.2 corrosion 3.2.1 electrolytic corrosion 3.2.2 applied voltages 3.2.3 connecting to different...
TRANSCRIPT
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
The Cost of Corrosion
Concrete International December 2004
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Electrolytic corrosion. When a metal is placed in water there is a tendency for it to dissolve (ionise) in the solution. Fe Fe++ + 2e-
where e- is the electron which remains in the metal. Positive metal ions are released into the solution and the process continues until sufficient negative charge has built up on the metal to stop the net flow.
-ve
Fe++
Electrode PotentialsMetal Electrode Potential
Magnesium -2.4
Aluminium -1.7
Zinc -0.76Chromium -0.65
Iron (ferrous) -0.44
Nickel -0.23Tin -0.14Lead -0.12Hydrogen (reference) 0.00
Copper (cupric) +0.34
Silver +0.80Gold +1.4
Current and exchange current
The current will depend exponentially on the difference between the potential and the rest potential:
where V is the Voltage across the anode and B1' is a constant for all samples
Similarly for the exchange current:
(1) e]B)/V[(V
I = I10a
aoa
(2) e]B)/ [(V
I = I10a
ao-aV
Anode current and exchange current
Anode current and exchange current
0.00E+00
2.00E-04
4.00E-04
6.00E-04
8.00E-04
1.00E-03
-1 -0.5 0
Voltage V
Cu
rren
t A
Ia+
Ia-
Notation for logarithms
Log(x) = Log to base 10
Ln(x) = Log to base e (natural log)
Thus Ln(x) = Ln(10)
Log(x)
The anode current
It may be seen that at voltages well above Va0
the exchange current is negligible and the voltage may be expressed by rearranging equation 1:
(3) I
I LogB + V = V0a
a10a
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Applied Potential to cause corrosion (reversed to stop it)
+
Fe++
Power
Supply
-
Current (electrons go the other way)
Cathodic protection
Cathodic Protection. Preparing the steel
(the cathode)
Cathodic protectionConductive paint anode (left)Titanium mesh anode (right)
Connection to rebar (left)Main junction box (right)
Bonding steel beams together (left), Casting in connection to the rebar (right)
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Zinc and Copper
Zinc anode system for reinforcement
protection
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
pH
pH = log(1/H+)where H+ is the number of grammes of hydrogen ions per litre.In pure water the following equilibrium reaction takes place:
H2O H+ + OH-
and there are 10-7 grammes of hydrogen ions per litre. Thus the pH of water is 7 and is defined as neutral. Acids have pH below 7 and alkalis (bases) have pH above 7. Concrete has a pH of 12.5.
Corrosion in pure water
The small amount which does take place is caused but the pH of water being 7, not infinite. i.e. there are 10-7 grammes of hydrogen ions per litre in neutral water. They are the product of the equilibrium of the reaction:
H2O H+ + OH-
in which the OH- is a hydroxyl ion which may combine with the iron ions in solution:
Fe++ + 2(OH)- Fe(OH)2 The product is ferrous hydroxide which is a green precipitate.
Anode and Cathode (Could be caused by applied voltage, different metals etc.)
+
Fe+++ 2(OH)- Fe(OH)2
-
CathodeAnode
e-
H2
H+
Anode and Cathode reaction
The reaction of the hydrogen ions with the electrons in the metal:
2H+ + 2e- H2
is known as the cathodic reaction and the dissolution of the metal ions:
Fe Fe++ + 2e-
is the anodic reaction.
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Corrosion with OxygenIf oxygen is present in the water it will react at the cathode:
2H2O + O2 + 4e- 4(OH)-
this uses up electrons at the cathode (increasing its potential) and provides hydroxyl ions to react with the iron ions in solution and thus greatly accelerates the corrosion. If there is a good supply of oxygen the final product is ferric hydroxide Fe(OH)3, this is common "red rust". If the air supply is limited, however, the product is Fe3O4 which is "black rust".
Oxygen
+
Fe++
-
CathodeAnode
e-
O2+ water 4(OH)-
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Corrosion in acids
Acids contain free positive hydrogen ions. Provided the metal has a potential below that of hydrogen the hydrogen ions will combine with the electrons in the metal to release hydrogen gas.
2H+ + 2e- H2 The metal ions will then combine with the acid in solution and the process will continue until either the metal or the acid is exhausted.
Acid corrosion
+
Fe++
-
CathodeAnode
e-
H2
H+
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Pitting
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Pourbaix diagram for steel
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Anode and cathode currents
Anode
CathodeCurrentIc
CurrentIa
Current Ix
Voltage V
STEEL +
CONCRETE -
The current will therefore depend on the resistance in the circuit.
This consists of :
• a. Surface resistance at the cathode
• b. Surface resistance at the anode
• c. Resistance in the solution
Solving for no applied voltageIf a cathodic process is initiated by any of the above processes (e.g. oxygen) its voltage may be expressed as:
If there is no external applied voltage the voltage is known as the rest potential Eo and the current flowing round the "loop"
is the corrosion current Icorr. Thus:
Thus subtracting from (3) and (4)
(4) LogB - V = V0c
c20c
I
I
(5) I
ILogB - V =
ILogB+ V=E
co
corr20c
0a
corr10a0
I
(6) I
I LogB- =
I
I LogB = EV
corr
c2
corr
a10
The linear approximation
but when x is close to 1: x-1 Ln(x)
Thus: (x-1) Log(x)
Ln(10)
Thus when Ia and Ic are close to Icorr
(7) 1-
I
I (10)LnB- = 1-
I
I (10)Ln
= EVcorr
c2
corr
a10
B
The Tafel Constants B1 and B2
and the Stern-Geary equationsWith the following definitions:
and
Equation (7) reduces to:
(8) (10)Ln)B+B(
BB = BConstant 21
21
(9) I
B = R resistanceon Polarisati
corr
p
(10) R
= I - I = ICurrent Externalp
0
cax
EV
Combination of Anode and Cathode Currents
-100-80-60-40-20
020406080
100
-600 -400 -200
Voltage mV
Cu
rren
t m
icro
A
Anode Current
Cathode Current
External Current(= Anode-Cathode)
Linear V-E0/Rp
Equivalent circuit for corroding surface
Resistance Rp
Diffusion Potential E0
Combination of Anode and Cathode Currents
-100-80-60-40-20
020406080
100
-600 -400 -200
Voltage mV
Cu
rren
t m
icro
A
Anode Current
Cathode Current
External Current(= Anode-Cathode)
Linear V-E0/Rp
Increased Anode Current (Ia0 increased)
-100
-80
-60
-40
-20
0
20
40
60
80
100
-600 -500 -400 -300 -200
Voltage mV
Cu
rren
t m
icro
A
Effect of increasing the anode current – Increased gradient indicating higher corrosion
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
NCE October 06
The corrosion Circuit
2Fe(OH)2Current
4e-
Electrons
2Fe++
4(OH)-
O2 and 2H2O
STEEL
CONCRETE
Cathodic reactionAnodic reaction
Pinholes in coating on bar – cathode will be 10 times larger than anode
Anode
Cathode
Offshore oil retaining structureAnode and cathode may be several metres apart
Air
Water Oil supplies oxygen to cathode
Cathode
Anode in splash zone
Large reinforced structure
Air (provides oxygen to cathode)
Cathode Anode
Two main differences in the equivalent circuit when in
concrete: 1. The circuit must pass through the concrete
which has a resistance. 2. The steel/concrete interface has a
capacitance. This is known as the "double layer capacitance" and is caused by charge build up at the interface.
Equivalent circuit of steel in concrete
The characteristics of the circuit1 If a voltage different from E0 is applied to it there will be a
high initial current through the capacitor but this will decay to zero. Thus in order to make a linear polarisation resistance measurement it is necessary either to:
•Wait about 30 second after applying the voltage. This has the disadvantage of causing possible changes to the corrosion process. or
•Apply a very slowly changing voltage. or
•Apply a pulse of voltage and make a measurement when it is switched off.
2 When measuring the polarisation resistance the concrete resistance will also be measured. Fortunately the capacitance has a very low resistance to alternating current so this may be used (50 - 100Hz) to measure the concrete resistance and it may then be subtracted.
Current decay
Experimental results for
linear polarisation
(high corrosion)
Experimental results for
linear polarisation
(low corrosion)
Equivalent circuit shown
with potentiostat
Potentiostat in use
Causing corrosion with an anodic voltage
Black rust from samplesTurns red when exposed to air
Circuit for resistance measurement
Measuring Resistivity
Potential Survey
Looking again at equation (5)
It may be seen that when comparing systems with similar cathode conditions (i.e. the same Ic0 and Vc0) as the rest potential Eo increases the log of the corrosion current Icorr decreases. This is the basis of a method of detecting corrosion called potential survey.
I
ILogB - V = E
0c
corr20c0
Rest potential
vs. corrosion current
Linear polarisation apparatus with guard ring
Polarization Resistance
Eo
Voltmeter
Ammeter
Switch D.C. Reference cell
Counter electrode
Working electrode
• Step 1: Measure open circuit potential, Eo
Eo + E
Ip
• Step 2: Close switch and apply small current• Step 3: Measure current, Ip, to produce small change in
voltage, E - 4 mV• Step 4: Increase current, and repeat measurement until E
-12 mV
Polarization Resistance, Rp
E
ip
Current/Area of Bar, ip, (µA/cm2)
Voltage Rp =E
ip
Corrosion Rate:
icorr = B
Rp
(µA/cm2)
B = 25 to 50 mV
From Faraday's Law:1 µA/cm2 = 0.012 mm/y
Guard-Electrode Method
Voltmeter
Ip
Guard ElectrodeAmmeter
VoltageFollower
Confine current so that affected area of bar is well defined
3.2 CORROSION• 3.2.1 Electrolytic corrosion• 3.2.2 Applied voltages• 3.2.3 Connecting to different metals• 3.2.4 Slow corrosion in pure water• 3.2.5 Oxygen• 3.2.6 Acids• 3.2.7 Pitting• 3.2.8 The effect of pH and potential• 2.3.9 Corrosion rates• 3.2.10 Corrosion of steel in concrete• 3.2.11 Corrosion Prevention
Corrosion Prevention
Coatings: This is the standard method (e.g. paint).Weathering Steels: Carbon steel with a 0.2% copper content forms a very stable oxide layer (in the absence of chlorides). It is therefore very durable, but equally ugly.Stainless Steels: are alloys of steel with some chromium and some other elements. Most stainless steels corrode to some extent..Cathodic protection: This method makes the metal cathodic (negative) relative to the solution and thus stops the anodic reaction.
Corrosion of stainless
steel
Sample panel of stainless
steel cladding
Corrosion Prevention
Coatings: This is the standard method (e.g. paint).Weathering Steels: Carbon steel with a 0.2% copper content forms a very stable oxide layer (in the absence of chlorides). It is therefore very durable, but equally ugly.Stainless Steels: are alloys of steel with some chromium and some other elements. Most stainless steels corrode to some extent..Cathodic protection: This method makes the metal cathodic (negative) relative to the solution and thus stops the anodic reaction.