Electrochemical Engineering: Topic 2: Corrosion Dr Rhodri Jervis Corrosion • What is corrosion? • The link between the electrochemical series and corrosion • The influence of pH on corrosion • Pourbaix diagrams • Measuring corrosion • Control and prevention of corrosion What is Corrosion? • How would you define corrosion? • The deterioration of materials through chemical processes • International standard definition of corrosion: “a physicochemical interaction between a metal and its environment which results in changes of the properties of the metal and which may often lead to impairment of the function of the metal, the environment, or the technical system of which these form a part.” Specifically, we’re interested here in the most significant form of corrosion industrially and economically: The electrochemical corrosion of metals to solution or their oxides Why is Corrosion Important? The World Corrosion Organisation state the following: • “The annual cost of corrosion world-wide exceeds the cost of all natural disasters • “The results of corrosion lead to pollution of the environment, the air and waters” • “Everything in our surroundings needs to be protected from corrosion” • “Corrosion is a ‘cancer’ on industry and needs to be solved” WCO Corrosion Awareness Day 2019 Why is Corrosion Important? Corrosion Costs: “At US $2.2 trillion, the annual cost of corrosion worldwide is over 3% of the world’s GDP. Yet, governments and industries pay little attention to corrosion except in high-risk areas like aircraft and pipelines. Now is the time for corrosion professionals to join together to educate industry, governments, and the public. Now is the time to work together to harmonize standards and practices around the world and to communicate and share corrosion mitigation technologies. Now is the time to make a major impact to protect the environment, preserve resources, and protect our fellow human beings. “Over the past 50 years, several national costs of corrosion studies have been conducted. Using different approaches, the studies all arrived at corrosion costs equivalent to about 3%–4% of each nation's gross domestic product (GDP). Using a 3.4% of global GDP (2013), the global cost of corrosion can then be estimated to be US$2.5 trillion. Using available corrosion control practices, it is estimated that savings of between 15% and 35% of the cost of corrosion could be realized, i.e., between US$375 and $875 billion annually on a global basis.” Now is the Time, George F. Hays, Director General World Corrosion Organization Koch, G. Trends in Oil and Gas Corrosion Research and Technologies. 2017. Web. Why is Corrosion Important? Safety There are many examples of industrial accidents caused, at least in part, by corrosion in the following areas (and more): • Aerospace • Chemical production • Oil and gas • Nuclear Read about some cases here: http://corrosiondoctors.org/Forms/Accidents.htm Aloha Incident Bhopal Disaster Noroco Refinery Explosion Davis-Besse Nuclear Reactor Why is corrosion important “Of the 137 major refinery accidents reported by EU countries to the EU’s eMARS database since 1984, around 20% indicated corrosion failure as an important contributing factor. This proportion of refinery accidents in eMARS with this profile has remained constant well into the 21st century” CorrosionβRelated Accidents in Petroleum Refineries, Wood, Vetere Arellano and Wijk, 2013 Electrochemistry and Corrosion ππ ↔ ππ ππ+ + ππ − • Consider the general reaction above • In the forward reaction, a metal loses electrons and forms a metal ion • Is this oxidation, or reduction? • Oxidation • What 3 things do we need for an electrochemical reaction to occur? • Electrodes (anode and cathode), electrolyte (ionic connection), electrical connection Electrochemistry and Corrosion • Recall the simple Daniel cell • What happens in galvanic (spontaneous) operation? • Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) • Zn metal tends to form Zn ions in solution, or in other words corrodes This can also be called a corrosion cell. How do we determine which metal corrodes (or gives up it’s electrons, or oxidises)? Galvanic Series/Reduction Potentials • The more negative standard reduction potential means that a metal has a greater tendency to oxidise (reverse of the reaction as written in these standard tables) • If two metals are in the same electrolyte and electrically connected, the one with the lowest standard reduction potential will corrode • These metals are some times called ‘base’, as opposed to noble metals that have more positive reduction potentials and are less likely to corrode Galvanic Series/Reduction Potentials • Metals such as iron tend to oxidise • What cathode reactions could be occurring (and what environmental conditions are required) to allow the oxidation of iron (or other metals at risk of corrosion)? Galvanic Series/Reduction Potentials • Hydrogen evolution reaction (HER) • Oxygen reduction reaction (ORR) (these reactions are highly important in electrochemistry and will come up in future lectures on fuel cells and electrolysers) Essential Elements of an Electrochemical Cell • Anode • Cathode • Electrolyte • Electrical connection • Separated reactions… Essential Elements of an Electrochemical Cell • Anode Oxidation of Fe πΉπΉπΉπΉ(π π ) → πΉπΉπΉπΉ 2+ + 2ππ − • Cathode ORR, HER • Electrolyte Water, moisture • Electrical connection The metal iron itself • Separated reactions… The anodic and cathodic reactions occur at different parts of the metal Cathode Reactions Oxygen reduction Reaction O2 + 4H+ + 4e- ↔ 2H2O O2 + 2H2O + 4e- ↔ 4OHHydrogen Evolution Reaction • 2H+ + 2e- ↔ H2 • 2H2O + 2e- ↔ 2OH- + H2 • ΔEo = -0.059 pH Eo = 1.23 V (pH = 0) Eo = 0.40 V (pH = 14) Eo = 0.00 V (pH = 0) Eo = -0.83 V (pH = 14) Which cathode reactions occur when? This depends on environmental and thermodynamic factors If a metal is relatively noble, and has a reduction potential above the hydrogen reaction, it is not possible for the HER to be the cathodic reaction. In the presence of oxygen, the ORR occurs at the cathode. Some strongly oxidising metals (reactive) such as Al, Zn, Mg can have a significant amount of hydrogen evolution occurring as they are reactive enough for the HER to be thermodynamically favourable. In the absence of oxygen, this is even more likely. The standard reduction potential series tells us how large the thermodynamic driving force could be for an electrochemical reaction. However, the kinetics of the reaction should be considered too. Galvanic Series/Reduction Potentials • Hydrogen evolution reaction (HER) • Oxygen reduction reaction (ORR) The ORR often gives a larger thermodynamic driving force, but the HER is more rapid (has faster kinetics) Other Cathode Reactions • Metal deposition: metal ions in solution can plate onto the surface • Metal ion reduction without plating: eg Fe(III) → Fe(II) • Reduction of carbonic acid H2CO3 + e- → H + CO3 • Reduction of chlorine: Cl2 + 2e-→ 2ClHowever, the ORR is the most common cathode reaction in normal conditions, except with highly reactive metals, in highly acidic solutions or where oxygen diffusion is limited pH Considerations How does pH affect corrosion, and why? What is pH? Recall that the standard reduction potential series is at standard conditions, meaning that the concentration of species in solution is set as 1 mol. The Nernst equation is used to calculate the change in thermodynamic potential with different concentration (or temperature). pH is the concentration of H+ ions, and so if they occur in the reduction reactions, the reversible potential of that half reaction will be altered in different pH pH ORR pH Dependence HER Dependence on pH General Expression of the Nernst Equation Taking the general equation for a half-cell reaction as: aA + mH+ + ze− = bB + H2O the Nernst equation becomes ORR and HER pH - diagram • As shown, the potential of the ORR/OER and HOR/HER vary with pH • The standard conditions are at pH 0 and so we get values of 1.23 and 0.00 V, respectively • Even though this changes with pH, the difference between them is always 1.23V (parallel lines) • This is the ‘breakdown voltage’ of water. • Above 1.23 V, water is broken down into oxygen, below 0.0 V into H2 Pourbaix Diagrams • This is also known as a Pourbaix diagram • It shows the thermodynamic stability of different species and their dependence on pH • In the previous case, this is Pourbaix diagram for water • The central zone is the stable region for water • Can think of the zones above and below this as being the ‘corrosion region’ of water • Pourbaix diagrams of other metals can give insight into the thermodynamics behind corrosion – i.e., are metals stable at particular regions Pourbaix Diagrams Pourbaix Diagram Rules • Each line represents a reversible reaction at thermodynamic equilibrium • Horizontal lines represent reactions that involve electron transfer (redox process) but not proton (or hydroxide) transfer • Eg ππππππ+ + ππππ− ↔ ππππ • Vertical lines represent acid-base reactions (proton or hydroxide) with no redox change • Eg ππππππππ+ + ππππππ πΆπΆ ↔ ππππππ πΆπΆππ + ππππ+ • Diagonal lines represent reactions where both electron (redox) and proton (or hydroxide) transfer occur • Eg ππππππππ πΆπΆππ + π―π―ππ πΆπΆ ↔ ππππππππ πΆπΆππ + ππππ+ + ππππ− Fe Pourbaix Diagram • The diagram is for Fe ion 1x10-6M • Thermodynamics only – no Kinetics • For specific Temperature and Pressure • Pure substance only, i.e. Fe not steel • Shows oxides to be thermodynamically stable, these are not necessarily passivating • Only anhydrous oxide species are shown and not all of the possible hydrated and non-stoichiometric thermodynamic species Fe2O3 = Fe(OH)3 Fe3O4 = FE(OH)2 Fe Pourbaix Diagram • Electrochemical equations can be written for the process of ‘crossing’ the line • Eg, line 1 would be: ππππππ+ + ππππ− ↔ ππππ • Line 2: ππππππ+ + ππ− ↔ ππππππ+ • Line 3: ππππππππ+ + ππππππ πΆπΆ ↔ ππππππ πΆπΆππ + ππππ+ • Line 4: ππππππππ+ + ππππππ πΆπΆ ↔ ππππππ πΆπΆππ + ππππ+ + ππππ− To evaluate the effect of a change of pH on the equilibrium position of one of these reactions, move horizontally. To evaluate the effect of applied electrode potential, move vertically Using Pourbaix to Predict Corrosion vs H2 evolution, Fe is unstable in acid and alkali Fe(s) + 2H+ Fe(s) + 2H2O Fe2+(aq) + H2 Fe3O4(s) + H2 (acid) (alkali) vs O2 reduction, Fe is also unstable in acid and alkali (as are most metals) 4Fe(s) + 3O2 + 12H+ 4Fe(s) + 3O2 4Fe3+ (aq) + 6H2O 2Fe2O3(s) (acid) (alkali) But at certain pH / potential regions the oxide will reduce corrosion http://chemwiki.ucdavis.edu More Pourbaix Diagrams • Gold’s Pourbaix diagram explains why it is the most immune substance known. It is immune in all regions in which cathodic reactions can take place. So gold never* corrodes in an aqueous environment. • Immunity of aluminium only occurs at lower potentials. Therefore, unless under conditions that cause it to passivate (pH 4-9), it is much more susceptible to corrosion than gold or zinc. * provided that the water is pure; that no ion complexes are present to provide a cathodic half cell reaction that occurs at a potential higher than +1.5 V(SHE). https://www.doitpoms.ac.uk/tlplib/pourbaix/printall.php Pros and Cons of Pourbaix Diagrams ο§ The diagrams provide an efficient pictorial summary of electron transfer, proton transfer and electron+proton transfer reactions which are favoured on a thermodynamic basis for a metal species in contact with a specific solution ο§ They should be used with caution: ο§ ο§ ο§ ο§ The pH of importance is that at the metal/electrolyte surface, which due to O2 reduction or H2 evolution might be very different to that of the bulk solution Although the diagram might indicate that a particular corrosion will occur spontaneously, this does not mean that significant corrosion will occur due to slow electron transfer kinetics or slow mass transport of ions Corrosion is complex, related to exact species in the electrolyte, metal/alloys, T, etc Thermodynamics shows the direction in which a reaction will tend; the rate and control of a corroding process is due to the Kinetics of the process which is revealed by studying the electrodics of the system Kinetics of Corrosion • The Nernst equation describes the thermodynamics of the reaction; i.e. when the system has attained equilibrium. In some systems this could take a long time (e.g. diamond-graphite). Pourbaix only shows the thermodynamics • The kinetics of the reaction are governed by: • The rate of electron transfer, or • The rate of diffusion of reactants and/or products to and from the metal surface. (More generally all mass transport processes) • The rate of electron transfer will depend on how well the metal surface catalyses the reaction. The magnitude of the exchange current density at equilibrium and the Tafel slope obtained from a voltamogram gives an indication of the reaction rate • If electron transfer is fast and the diffusion of reactants and/or products to and from the metal surface is slow, then the rate is controlled by diffusion. e.g. mass transport inhibited by an oxide layer on the surface of the metal Passivation ο§ If a solid oxide is formed on the metal, then this can inhibit O2 diffusion to the metal surface. This is called ‘passivation’. The effectiveness of the passive layer depends on: ο§ Its porosity to O2 transfer and Mz+ transfer ο§ Whether the oxide layer stays on or falls off (mechanical stability) ο§ Whether other ions are present in the passive layer ο§ Passive layers can form and be dissolved in relation to the potential of the metal and the local pH of the electrolyte at the metal/solution surface. ο§ ο§ The local pH can change if the reaction involves H+ or OH- ions The Pilling–Bedworth ratio is the ratio of the volume of the crystallographic unit cell of the metal oxide to that of the metal from which the oxide is formed. ο§ It can be used to predict if a passivation layer will form in dry air ο§ RPB < 1 layer too thin and cracks; RPB > 2 layer has internal strain and falls off (Fe); 1 < RPB < 2: the oxide coating is passivating. There are however numerous exceptions. Passivating metals: Al, Ti, Chromium containing steel Measuring Corrosion • For corrosion to be possible it has to be thermodynamically favourable – i.e, a spontaneous reaction, a reduction in Gibbs free energy of the system • For corrosion to be a significant problem, the kinetics have to be relatively fast – i.e., the rate of corrosion has to be such that nonnegligible degradation of a metal occurs over the timescale of years • However, most corrosion is relatively slow, and so methods such as measuring mass loss of a metal might not work for practical experimentation – electrochemistry can be used to measure the corrosion rate in less time A reminder of the Tafel approximation of the Butler Volmer reaction describing electrode kinetics: Reaction Kinetics Tafel relationship π±π± = π±π±ππ πΆπΆπ¨π¨ ππππ −πΆπΆπͺπͺ ππππ ππππ π©π© πΌπΌ − ππππ π©π© πΌπΌ πΉπΉπΉπΉ πΉπΉπΉπΉ • This equation can be re-written such that it takes the form of the Tafel relationship which was observed experimentally: η = a + b log (J) high +η: πΆπΆπ¨π¨ ππππ πΌπΌ πΉπΉπΉπΉ −πΆπΆπͺπͺ ππππ ππππ π©π© πΌπΌ πΉπΉπΉπΉ π±π± = π±π±ππ ππππ π©π© high -η: −π±π± = −π±π±ππ Tafel Plot Region ln(i) πΆπΆπ¨π¨ πππΌπΌ + π₯π₯π₯π₯ π±π±ππ π₯π₯π₯π₯ π±π± = πΉπΉπΉπΉ −πΆπΆπͺπͺ πππΌπΌ + π₯π₯π₯π₯ π±π±ππ π₯π₯π₯π₯ π±π± = πΉπΉπΉπΉ Tafel Plot Region ln(i0) -η [14] Katherine Holt - UCL +η Using Electrochemistry in Corrosion • Because corrosion occurs via electrochemical reactions, electrochemical techniques are ideal for the study of the corrosion processes. In electrochemical studies a metal sample a few cm² in surface area is used to model the metal in a corroding system. The metal sample is immersed in a solution typical of the metal's environment in the system being studied. Additional electrodes are immersed in the solution, and all the electrodes are connected to a device called a potentiostat. A potentiostat allows you to change the potential of the metal sample in a controlled manner. • The DC Corrosion standard techniques use the potentiostat to perturb the equilibrium corrosion process. When the potential of a metal sample in solution is forced away from Eoc, we call it polarizing the sample. The response (current or voltage) of the metal sample is measured as it is polarized. The response is used to develop a model of the sample's corrosion behaviour. Both controlled potential (potentiostatic) and controlled current (galvanostatic) polarization are useful. When the polarization is done potentiostatically, current is measured, and when it is done galvanostatically, potential is measured. LSV In Corrosion Analysis log • Linear sweep voltammetry can be used to sweep the potential of an electrode (metal of interest) and measure the current • A typical corrosion experiment is shown • Around the equilibrium potential (Ecorr), reduction and oxidation (corrosion) processes occur • There is then a passivation region where foration of passivating oxides means the current lowers to zero • Eventually, the potential is high enough to break down the passive layer and oxidation can again occur (this time at an accelerated rate due to the high overpotential) LSV In Corrosion Analysis log As with the Tafel analysis, a plot of the log of the current density is used to extract kinetic parameters from the data LSV In Corrosion Analysis Tafel Analysis providing π¬π¬ππππππππ and ππππππππππ Stainles s Steel in 1 M H2SO4 + 2ppm F- 0 1.000x10 -1 1.000x10 -2 1.000x10 -3 i/A 1.000x10 1.000x10 -4 1.000x10 -5 1.000x10 -6 1.000x10 -7 1.000x10 -8 -0.75 Corros ion rate analy s is -0.50 -0.25 0 0.25 0.50 E/V • Plot log absolute current vs potential (i.e. the magnitude of the current NOT the ‘direction’) • Calculate gradient of slope at +/- 50mV from the minimum point • Intersection of lines gives Ecorr and icorr values 0.75 1.00 1.25 1.50 LSV In Corrosion Analysis Tafel Analysis providing π¬π¬ππππππππ and ππππππππππ Stainles s Steel in 1 M H2SO4 + 2ppm F- 0 1.000x10 -1 1.000x10 -2 1.000x10 -3 i/A 1.000x10 1.000x10 -4 1.000x10 -5 1.000x10 -6 1.000x10 -7 1.000x10 -8 • Plot log absolute current vs potential (i.e. the magnitude of the current NOT the ‘direction’) • Calculate gradient of slope at +/- 50mV from the minimum point • Intersection of lines gives Ecorr and icorr values Corros ion rate analy s is icorr -0.75 Ecorr -0.50 -0.25 0 0.25 0.50 E/V 0.75 1.00 1.25 1.50 LSV In Corrosion Analysis Alternative Corrosion Plot – Voltage vs. log Current (A/m2) Corrosion plots are often plotted with log current density on the x axis The linear region of the Tafel plots form tangents and cross at a point that defines the corrosion current (i ) and the corrosion potential (e ) LSV In Corrosion Analysis The following figure illustrates this process. The vertical axis is potential and the horizontal axis is the logarithm of absolute current. The theoretical current for the anodic and cathodic reactions are shown as straight lines. The curved line is the sum of the anodic and cathodic currents. When you sweep the potential of the metal, you measure the current. The sharp point in the curve results from the use of a logarithmic axis. It is actually the point where the current gets very small prior to changing sign. Gamry.com Measuring corrosion rate from ππππππππππ Remember: Faraday’s Law: the amount of deposition or loss of a substance caused by an electrochemical reaction is proportional to the current passed Corrosion rate (CR): icorr is the corrosion current (A/m2) M is atomic mass of Fe = 55.85*10-3 kgmol-1 ρ is density of Fe = 7876 kgm-3 n is the number of electrons transferred F is Faraday constant= 96.485 C/mol ππππππππππ . π΄π΄ πͺπͺπͺπͺ = ππππππ Exercise: Perform dimensional analysis to determine the standard units of corrosion rate Corrosion prevention For corrosion to occur, both the anodic and cathodic processes must happen There are therefore several methods used to prevent or slow corrosion Corrosion prevention: Coatings • The most simple method is to coat the surface with something that acts as a barrier to moisture and oxygen • Paints, polymers, rubbers etc • Breaching of the coating could easily occur, leaving the metal vulnerable to corrosion Prevention: Sacrificial Coatings • An imposed negative charge on the metal drives the reversible equilibrium towards reduction, and so oxidation is less likely to happen. However, this is often not practical • One way to supply this negative charge is to coat with a sacrificial coating • The is where a more active metal is coated and goes becomes the anodic reaction in place of the metal that is being protected • This means the sacrificial metal dissolves preferentially • Zn is a common coating method – this is called galvanization Prevention: Cathodic Protection • A sacrificial anode can also be used to maintain a continual negative charge on the metal to be protects • A separate piece of metal is connected electrically to the metal of interest. No current needs to be supplied – the corrosion of the sacrificial anode supplies the charge Prevention: Anodising • A positive potential is applied to the metal to force a passivation coating • Increases the natural oxide layer thickness on a metal • Controlled growth of good quality oxide layer prevents it from being easily broken down • Can cause the metal to become insulating (can be problematic for some applications) • Commonly used with aluminium Prevention: Alloying • Alloying with other metals can make a metal more resistant to corrosion • Stainless steel is iron containing chromium, manganese, silicon, carbon and sometimes nickel and molybdenum • These elements react with moisture/water and air to form a very thin (a few atomic layers) stable film formed of corrosion products such as oxides and hydroxide • Chromium is the most important element in this – usually at 10 percent chromium or more • The film acts as a barrier layer to moisture and oxygen • The film is so thin that, although it’s technically corrosion it’s not possible to see and thus the steel appears ‘stainless’ Summary • Corrosion is a major industrial issue • Thermodynamic and kinetic aspects need to be taken into consideration when assessing the likelihood and damage of corrosion • Its root causes are electrochemical, and we can use electrochemical methods to study and even prevent corrosion Example Problems • Steel corrodes via the following electrochemical half reaction: • πΉπΉπΉπΉ(π π ) → πΉπΉπΉπΉ 2+ + 2ππ − • A corrosion current is measured and shown in the plot on the right • Calculate the corrosion rate in mm/year V Log (mA/cm2) Corrosion rate (CR): ππππππππππ . π΄π΄ πͺπͺπͺπͺ = ππππππ icorr is the corrosion current (A/m2) M is atomic mass of Fe = 55.85*10-3 kgmol-1 ρ is density of Fe = 7876 kgm-3 n is the number of electrons transferred F is Faraday constant= 96.485 C/mol
