# 3.2 CORROSION Electrolytic corrosion Applied voltages

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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 Corrosion of steel in concrete 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 Corrosion of steel in concrete 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 Potentials Metal Electrode Potential Magnesium -2.4
Aluminium -1.7 Zinc -0.76 Chromium -0.65 Iron (ferrous) -0.44 Nickel -0.23 Tin -0.14 Lead -0.12 Hydrogen (reference) 0.00 Copper (cupric) +0.34 Silver +0.80 Gold +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:

Anode current and exchange current

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:

3.2 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 Corrosion of steel in concrete Corrosion Prevention

Applied Potential to cause corrosion (reversed to stop it)
+ - Power Supply Current (electrons go the other way) Fe++

Cathodic protection

Cathodic Protection. Preparing the steel (the cathode)

Cathodic protection Conductive 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 Corrosion of steel in concrete 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 Corrosion of steel in concrete 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.)
+ - Anode Cathode H2 H+ e- Fe+++ 2(OH)-  Fe(OH)2

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 Corrosion of steel in concrete Corrosion Prevention

Corrosion with Oxygen If 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 + - Anode Cathode O2+ water 4(OH)- e- Fe++

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 Corrosion of steel in concrete 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 + - Anode Cathode H2 H+ e- Fe++

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 Corrosion of steel in concrete 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 Corrosion of steel in concrete 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 Corrosion of steel in concrete Corrosion Prevention

Anode and cathode currents
Ic Ia Current Ix Voltage V STEEL + CONCRETE -

The current will therefore depend on the resistance in the circuit
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 voltage
If 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)

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

The Tafel Constants B1 and B2 and the Stern-Geary equations
With the following definitions: and Equation (7) reduces to:

Equivalent circuit for corroding surface
Resistance Rp Diffusion Potential E0

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 Corrosion of steel in concrete Corrosion Prevention

NCE October 06

The corrosion Circuit CONCRETE 4e- STEEL O2 and 2H2O Current 2Fe(OH)2
Electrons 2Fe++ 4(OH)- O2 and 2H2O STEEL CONCRETE Cathodic reaction Anodic reaction

Pinholes in coating on bar – cathode will be 10 times larger than anode

Offshore oil retaining structure Anode and cathode may be several metres apart
Air Anode in splash zone Water Oil supplies oxygen to cathode Cathode

Large reinforced structure
Cathode Anode Air (provides oxygen to cathode)

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 circuit
1 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 ( Hz) 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 samples Turns 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.

Rest potential vs. corrosion current

Linear polarisation apparatus with guard ring

Polarization Resistance
Step 1: Measure open circuit potential, Eo Voltmeter Ammeter Switch D.C. Eo Reference cell Counter electrode Working electrode

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 Eo +  E Ip

Polarization Resistance, Rp
Corrosion Rate: icorr = B Rp (µA/cm2) Rp = E  ip Voltage E ip B = 25 to 50 mV From Faraday's Law: 1 µA/cm2 = mm/y Current/Area of Bar, ip, (µA/cm2)

Guard-Electrode Method
Confine current so that affected area of bar is well defined Voltmeter Voltage Follower Ip Ammeter Guard Electrode

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 Corrosion of steel in concrete 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.