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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.

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Presentation on theme: "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."— Presentation transcript:

1 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

2 The Cost of Corrosion Concrete International December 2004

3 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

4 Electrolytic corrosion. When a metal is placed in water there is a tendency for it to dissolve (ionise) in the solution. Fe Fe e - 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 ++

5 Electrode Potentials MetalElectrode 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

6 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 B 1 ' is a constant for all samples Similarly for the exchange current:

7 Anode current and exchange current

8 Notation for logarithms Log(x) = Log to base 10 Ln(x) = Log to base e (natural log) Thus Ln(x) = Ln(10) Log(x)

9 The anode current It may be seen that at voltages well above V a0 the exchange current is negligible and the voltage may be expressed by rearranging equation 1:

10 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

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

12 Cathodic protection

13 Cathodic Protection. Preparing the steel (the cathode)

14 Cathodic protection Conductive paint anode (left) Titanium mesh anode (right)

15 Connection to rebar (left) Main junction box (right)

16 Bonding steel beams together (left), Casting in connection to the rebar (right)

17 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

18 Zinc and Copper

19 Zinc anode system for reinforcement protection

20 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

21 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: H 2 O H + + OH - and there are 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.

22 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 grammes of hydrogen ions per litre in neutral water. They are the product of the equilibrium of the reaction: H 2 O H + + OH - in which the OH - is a hydroxyl ion which may combine with the iron ions in solution: Fe (OH) - Fe(OH) 2 The product is ferrous hydroxide which is a green precipitate.

23 Anode and Cathode (Could be caused by applied voltage, different metals etc.) + Fe (OH)- Fe(OH)2 - CathodeAnode e-e- H2H2 H+H+

24 Anode and Cathode reaction The reaction of the hydrogen ions with the electrons in the metal: 2H + + 2e - H 2 is known as the cathodic reaction and the dissolution of the metal ions: Fe Fe e - is the anodic reaction.

25 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

26 Corrosion with Oxygen If oxygen is present in the water it will react at the cathode: 2H 2 O + O 2 + 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 Fe 3 O 4 which is "black rust".

27 Oxygen + Fe ++ - CathodeAnode e-e- O 2 + water 4(OH)-

28 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

29 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 - H 2 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.

30 Acid corrosion + Fe ++ - CathodeAnode e-e- H2H2 H+H+

31 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

32 Pitting

33 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

34 Pourbaix diagram for steel

35 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

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

37 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

38 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 E o and the current flowing round the "loop" is the corrosion current I corr. Thus: Thus subtracting from (3) and (4)

39 The linear approximation but when x is close to 1: x-1 Ln(x) Thus: (x-1) Log(x) Ln(10) Thus when I a and I c are close to I corr

40 The Tafel Constants B 1 and B 2 and the Stern-Geary equations With the following definitions: and Equation (7) reduces to:

41

42 Equivalent circuit for corroding surface Resistance R p Diffusion Potential E 0

43 Effect of increasing the anode current – Increased gradient indicating higher corrosion

44 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

45 NCE October 06

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

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

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

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

50 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.

51 Equivalent circuit of steel in concrete

52 The characteristics of the circuit 1 If a voltage different from E 0 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. 2When 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.

53 Current decay

54 Experimental results for linear polarisation (high corrosion)

55 Experimental results for linear polarisation (low corrosion)

56 Equivalent circuit shown with potentiostat

57 Potentiostat in use

58 Causing corrosion with an anodic voltage Black rust from samples Turns red when exposed to air

59 Circuit for resistance measurement

60 Measuring Resistivity

61 Potential Survey Looking again at equation (5) It may be seen that when comparing systems with similar cathode conditions (i.e. the same I c0 and V c0 ) as the rest potential E o increases the log of the corrosion current I corr decreases. This is the basis of a method of detecting corrosion called potential survey.

62 Rest potential vs. corrosion current

63 Linear polarisation apparatus with guard ring

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

65 E o + E IpIp Step 2: Close switch and apply small current Step 3: Measure current, I p, to produce small change in voltage, E - 4 mV Step 4: Increase current, and repeat measurement until E -12 mV

66 Polarization Resistance, R p E i p Current/Area of Bar, i p, (µA/cm 2 ) Voltage R p = E i p Corrosion Rate: i corr = B RpRp (µA/cm 2 ) B = 25 to 50 mV From Faraday's Law: 1 µA/cm 2 = mm/y

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

68 3.2 CORROSION Electrolytic corrosion Applied voltages Connecting to different metals Slow corrosion in pure water Oxygen Acids Pitting The effect of pH and potential Corrosion rates Corrosion of steel in concrete Corrosion Prevention

69 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.

70 Corrosion of stainless steel

71 Sample panel of stainless steel cladding

72 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.


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