SR_Chemistry_Notes on Corrosion_Part II.pdf

Corrosion What is Corrsion? “The phenomenon of deterioration and destruction of matter by unwanted, unintentional attack of the environment leading to loss of matter starting at its surface is called corrosion”. Corrosion is a natural process,...

Corrosion What is Corrsion? “The phenomenon of deterioration and destruction of matter by unwanted, unintentional attack of the environment leading to loss of matter starting at its surface is called corrosion”. Corrosion is a natural process, which converts a refined metal to a more stable form, such as its oxide, hydroxide, or sulfide. It is the gradual destruction of materials (usually metals) by chemical and/or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to controlling and stopping corrosion. Corrosion can be defined as the degradation of a material due to a reaction with its environment. Examples i) Rusting of iron – when iron is exposed to the atmospheric conditions, a layer of reddish scale and powder of Fe3O4 is formed. ii) Formation of green film of basic carbonate- [CuCO3 + Cu(OH)2] on the surface of copper when exposed to moist air containing CO2. 1  Consequences of Corrosion The consequences of corrosion are many and varied and the effects of these on the safe, reliable and efficient operation of equipment or structures are often more serious than the simple loss of a mass of metal. Failures of various kinds and the need for expensive replacements may occur even though the amount of metal destroyed is quite small. Some of the major harmful effects of corrosion can be summarised as follows: Reduction of metal thickness leading to loss of mechanical strength and structural failure or breakdown. When the metal is lost in localised zones so as to give a cracklike structure, very considerable weakening may result from quite a small amount of metal loss. Hazards or injuries to people arising from structural failure or breakdown (e.g. bridges, cars, aircraft). Loss of time in availability of profile-making industrial equipment. Reduced value of goods due to deterioration of appearance. Contamination of fluids in vessels and pipes (e.g. beer goes cloudy when small quantities of heavy metals are released by corrosion). Perforation of vessels and pipes allowing escape of their contents and possible harm to the surroundings. For example a leaky domestic radiator can cause expensive damage to carpets and decorations, while corrosive sea water may enter the boilers of a power station if the condenser tubes perforate. Loss of technically important surface properties of a metallic component. These could include frictional and bearing properties, ease of fluid flow over a pipe surface, electrical conductivity of contacts, surface reflectivity or heat transfer across a surface. Mechanical damage to valves, pumps, etc, or blockage of pipes by solid corrosion products. Added complexity and expense of equipment which needs to be designed to withstand a certain amount of corrosion, and to allow corroded components to be conveniently replaced. The losses incurred are very huge and it is estimated that the losses due to corrosion are approximately 2 to 2.5 billion dollars per annum all over the world. 2  Types of Corrosion Corrosion Dry or Chemical Corrosion Wet or Electrochemical corrosion ❑ Dry Corrosion The corrosions that involve the direct attack of the atmospheric gases, generally in the absence of moisture, (i.e., conducting medium) is called as dry corrosion or chemical corrosion. Atmospheric gases include oxygen, halogens, oxides of sulphur, nitrogen, hydrogen sulphides etc. Dry Corrosion Oxidation Corrosion Corrosion by other gases Liquid metal corrosion 3  Dry Corrosion Oxidation Corrosion This is brought about by the direct action of oxygen on the metal surface in the absence of moisture. The oxygen atoms of the air are held close to the surface by means of weak Vander waals forces. Over a period of time, these forces results in the formation of weak bonds converting the metal into its corresponding metal oxide. The phenomenon is known as chemisorption. Oxidation occurs first at the surface of the metal to form metal oxides as per the following reaction: 4M 4 Mn+ + 4 ne Oxidation n O2 + 4 ne n O2- Reduction 4 M + n O2 4 Mn+ + n O2- 4  Dry Corrosion Mechanism Oxygen anions Oxidation Occurs at air- oxide interface Oxidation Occurs at the metal-oxide interface Metal Cations ❑ Initially the surface of metal undergoes oxidation and the resulting metal oxide scale forms a barrier which restricts further oxidation. The extent of corrosion depends upon the nature of metal oxide. ❑ The nature of the oxide film formed plays an important role on the surface of the metal as it may be stable, unstable, volatile and porous. 5  Dry Corrosion Stable Metal Oxides It is impervious in nature and forms a protective coating, thus preventing the further oxidation of metals. Examples: oxide films of Al, Pb, Cu, Sn etc Exposed Surface Stable Metal Oxide + O2 Metal Metal No further decomposition of air 6  Dry Corrosion Exposed Surface Unstable Metal Oxide Unstable Metal Oxides + O2 Metal Oxide This type of oxide layer decomposes back to metal and oxygen as soon as Metal Metal Metal of air Decomposes it is formed. So no oxidation corrosion is observed for the metals. Examples: Oxides of Ag and Au. Fresh Surface Volatile Metal Oxides Exposed Surface Volatile Metal Oxide exposed for further attack The metal oxide film volatizes as soon as it is formed, thereby leaving the + O2 Metal Oxide underlying metal exposed to the environment for further attack. Thus, rapid Metal Metal Metal of air Volatizes corrosion takes place. Example: Molybdenum oxide (MoO3). Porous Metal Oxides Exposed Surface Porous Metal Oxide The oxide layers have pores and cracks and allow the atmospheric oxygen + O2 Further attack through to access the underlying layers of metals. This results in unobstructed and Metal Metal of air pores/cracks continues rapid corrosion of the metal. The process continues until the entire metal is converted into its oxides. Examples: Oxides of Li, Mg, Na etc. 7  Dry Corrosion Pilling Bedworth Rule ▪ To express the extent of protection given by the corrosion layer to the underlying metal Pilling Bedworth rule was postulated. ▪ It is expressed in terms of specific volume ratio. According to this, “an oxide product is protective or non-porous, if the volume of oxide is at least as great as the volume of metal from which it is formed”. On the other hand, if the volume of oxide formed is less than the volume of the metal, the oxide layer is porous and non-protective. Thus smaller is the specific volume ratio (Volume of metal oxide/Volume of the metal), greater is the oxidation corrosion. Specific Volume ratio= 𝐕𝐨𝐥𝐮𝐦𝐞 𝐨𝐟 𝐦𝐞𝐭𝐚𝐥 𝐨𝐱𝐢𝐝𝐞 𝐥𝐚𝐲𝐞𝐫 𝐕𝐨𝐥𝐮𝐦𝐞 𝐨𝐟 𝐩𝐚𝐫𝐞𝐧𝐭 𝐦𝐞𝐭𝐚𝐥 ▪ Smaller the specific volume ratio, greater is the oxidation corrosion Eg. The specific volume ratio of W, Cr, and Ni are 3.6, 2.0 and 1.6 respectively. Consequently the rate of corrosion is least in Tungsten(W) ❑ If the volume of the corrosion film formed is more than the underlying metal, it is strongly adherent, non-porous and does not allow the penetration of corrosive gases. No further corrosion. ❑ If the volume of the corrosion film formed is less than the underlying metal, it forms pores/cracks and allow the penetration of corrosive gases leading to corrosion of the underlying metal. 8  Dry Corrosion Corrosion by other gases This type of corrosion takes place by the chemical affinity of gases such as SO2, CO2, Cl2, H2S, and F2 etc. The corrosive effect depends on the chemical affinity between the metal and gas, and the formation of a protective or non-protective layer on the metal surface. Example, AgCl forms the protective films. SnCl4 forms a volatile product, while attack of Fe by H2S gas forms a porous FeS film. Liquid metal corrosion Liquid metal corrosion is a physical or physical-chemical process that occurs when liquid metal interacts with solid materials, removing the solid's constituents as solutes or forming compounds. The process can be broken down into three steps: ▪ Transport in the solid ▪ Dissolution of the steel constituents into the liquid ▪ Transport of the corrosion products and impurities in the liquid For example, liquid mercury dissolves most metals by forming amalgams, thereby corroding them. Also occurs in devices used for nuclear power plant in cooling system. Coolant (liquid sodium metal) leads to corrosion of cadmium in nuclear plant. 9  Corrosion Penetration Rate Corrosion Penetration Rate (CPR) Corrosion penetration rate is defined in three ways ❖ The speed at which any metal in a specific environment deteriorates due to chemical reaction in the metal when it is exposed to a corrosive environment. ❖ The amount of corrosion loss per year in thickness ❖ The speed at which corrosion spreads to the inner portions of a material. The CPR is calculated as follows Units of CPR 𝜅. 𝑊 CPR in mpy (mil per year) 𝐶𝑃𝑅 = 𝜌. 𝐴. 𝑡 1 mil = 0.001 inch k = 534; W → mg; ρ → g cm-3; t → hours Where, A → inch2 k is constant W is the total weight loss CPR in mmpy (milimeter per year) ρ is the density of metal in g.cm-3 A is the area of the exposed metal k = 87.6; W → mg; ρ → g cm-3; t → hours t is the time taken for the loss of the metal A → cm2 10  Problem on CPR ▪ A piece of corroded steel plate was found in a submerged ocean vessel. It was estimated that the original area of the plate was 10 inch2 and approximately 2.6 kg has corroded away during the submersion. Assuming a CPR of 200 mpy for this alloy in sea waters, calculate the time of submersion in years. The density of steel is 7.9 g.cm-3 k = 534; W = 2.6 kg; ρ = 7.9 g.cm-3; A = 10 inch2; t=? 𝜅. 𝑊 𝐶𝑃𝑅 = 𝜌. 𝐴. 𝑡 ▪ A steel sheet of area 100 inch2 is exposed to air near the ocean. After 1 year period it was found to experience a weight loss of 485 g due to corrosion and density of steel metal is 7.9 g.cm-3. Calculate the CPR in mpy. k = 534; W = 0.485 kg; ρ = 7.9 g.cm-3; A = 100 inch2; t = 1 year 𝜅. 𝑊 𝐶𝑃𝑅 = 𝜌. 𝐴. 𝑡 11  Wet Corrosion This type of corrosion occurs when a conducting liquid is in contact with metal or when two dissimilar metals or alloys are either immersed or dipped partially in a solution. ❖ It involves the formation of two areas of different potentials in contact with a conducting liquid. One is named as anodic area where oxidation reaction takes place, the other is referred to as a cathodic area involving reduction. ❖ The metal at anodic area is destroyed either by dissolving or by forming a combined state, such as oxides. Hence corrosion always occurs at anodic areas. ❖ At cathode, the dissolved constituents gain the electrons forming non-metallic ions. The metallic ions and non-metallic ions diffuse towards each other forming product somewhere between anode and cathode. 12  Wet Corrosion Anode Electro chemical corrosion involves flow of electric current between anodic and cathodic areas. At anode, dissolution of metal takes place forming corresponding metallic ions. 4M 4 Mn+ + 4 ne Oxidation Cathode On the other hand, at cathode, consumption of electrons takes place either by H+ H+ H+ Acidic Solution H+ i) Evolution of hydrogen type (Electrolyte) Diffusion of Ferrous ii) Absorption of oxygen type Ions into the electrolyte 2H+ + 2 e H2 Cathodic Reaction Fe Fe2+ + 2 e Fe Fe2+ + 2 e Evolution of hydrogen type e– This type of corrosion occurs if the conducting medium is acidic in nature. For example, Iron Small Cathodic Area Anodic Area Large Anodic Area Large dissolves and forms ferrous ions with the liberation of electrons. These electrons flow from anode to cathode, where H+ ions are eliminated as hydrogen gas. Iron 2H+ + 2 e H2 The electrons released at anode flow through the metal from anode to cathode, where as H+ ions of acidic solution take up these electrons and eliminated as hydrogen gas (Reduction) Fe Fe2+ + 2 e At Anode dissolution of iron to ferrous ion takes place with the liberation of electrons (Oxidation) 13 Fe + 2H+ Fe2+ + H2 ▪ This type of corrosion causes “displacement of hydrogen ions from the acidic solution by metal ions. ▪ In hydrogen evolution type corrosion, the anodic areas are large and cathodic areas are small.  Wet Corrosion Absorption of Oxygen A cathodic reaction can be absorption of oxygen, if the conducting liquid is neutral or aqueous and sufficiently aerated. Some cracks developed in iron oxide film cause this type of corrosion. The surface of iron is always coated with a thin oxide film. The crack developed will create an anodic area on the surface while the well coated metal parts act as cathode. The anodic areas are small and the cathodic areas are large. Corrosion occurs at the anode and rust occurs in between anode and cathodic areas. When the amount of oxygen increases corrosion is accelerated. ▪ This type of corrosion takes place in basic or neutral medium in presence of oxygen. ▪ For example, rusting of iron in neutral or basic aqueous solution of electrolyte in presence of atmospheric oxygen. Aqueous Neutral Solution of Electrolyte ▪ Usually the surface of iron is coated with a thin film of iron oxide. Rust ▪ If the film develops cracks, anodic areas are created on the surface and the rest of the 𝟏 𝑶 + 𝟐𝒆 + 𝑯𝟐 𝑶 𝟐𝑶𝑯− 𝟏 𝑶 + 𝟐𝒆 + 𝑯𝟐 𝑶 𝟐𝑶𝑯− metal surface acts as cathodes. 𝟐 𝟐 𝟐 𝟐 Oxide Film ▪ It shows that anodic areas are small and the cathodic areas are large. Cathode e– e– Cathode (Large) (Large) Flow of Electron Flow of Electron Anode Fe Fe2+ + 2 e The Fe2+ ions formed at anode, and OH- ions formed at cathode, diffuse towards Small Anodic area caused by cracks in 𝟏 the oxide film where corrosion occurs each other forming Fe(OH)2 i.e., Cathode 𝑶 + 𝟐𝒆 + 𝑯𝟐 𝑶 𝟐𝑶𝑯− Fe Fe2+ + 2 e 𝟐 𝟐 Fe2+ + 2 OH- Fe(OH)2 Iron 𝟏 𝑶 + 𝑭𝒆 + 𝑯𝟐 𝑶 𝑭𝒆𝟐+ + 𝟐𝑶𝑯− 𝟐 𝟐 𝟒𝑭𝒆(𝑶𝑯)𝟐 + 𝑶𝟐 + 𝟐𝑯𝟐 𝑶 𝟒𝑭𝒆(𝑶𝑯)𝟑 𝟒𝑭𝒆(𝑶𝑯)𝟑 Oxidation 𝑭𝒆𝟐 𝑶𝟑. 𝒙𝑯𝟐 𝑶 Rust (hydrated ferric oxide) If enough oxygen is present, ferrous hydroxide is easily oxidized to ferric hydroxide and then to hydrated ferric oxide which is known as rust. 14  Difference between chemical Corrosion and electrochemical corrosion Chemical Corrosion Electrochemical Corrosion It takes place in dry condition It takes place in wet condition such as in the presence of electrolytes. It involves the direct chemical attack of environment of the It involves the formation of large number of galvanic cells. metal. It takes place on homogeneous and heterogeneous surfaces. It takes place on heterogeneous surfaces only. Corrosion product accumulates at the same place where Corrosion product accumulates at cathode, but corrosion takes corrosion is taking place. place at anode. Uniform corrosion takes place. Non – Uniform corrosion takes place 15  Galvanic Series The galvanic series (or electropotential series) determines the nobility of metals and semi- metals. When two metals are submerged in an electrolyte, while also electrically connected by some external conductor, the less noble (base) will experience galvanic corrosion. The rate of corrosion is determined by the electrolyte, the difference in nobility, and the relative areas of the anode and cathode exposed to the electrolyte. The difference can be measured as a difference in voltage potential: the less noble metal is the one with a lower (that is, more negative) electrode potential than the nobler one, and will function as the anode (electron or anion attractor) within the electrolyte device functioning as described above (a galvanic cell). Galvanic reaction is the principle upon which batteries are based. Figure shows the galvanic series measured in seawater for some common metals and alloys. When two metals are further apart in the list (e.g. a larger difference between the two numbers), the driving force for galvanic corrosion is increased. The most anodic (active) metals are at the top and most cathodic (noble) at the bottom. Both solid and hollow bars are shown for the stainless steels. The hollow bars represent actively corroding stainless steel, which has a different potential then passive (not corroding) stainless steel. In most applications, where dissimilar metals are combined, the passive (solid) bar should be used to determine the position of the stainless steel. For example, if zinc (think galvanized steel) which is an active material and near the top of the list and stainless steel, a noble metal and near the bottom of the list were in direct contact and in the presence of an electrolyte (water), galvanic corrosion will occur if they are regularly exposed to an electrolyte. 16  Different Types of Electrochemical Corrosion Galvanic Corrosion or Bimetallic Corrosion When two dissimilar metals are electrically connected and exposed to an electrolyte, the metals higher in electrochemical series have a tendency of forming anode and undergo corrosion. For example, when zinc and copper are electrically connected either in acidic solutions or in their respective salt solution, zinc being more anodic by virtue of its position in electro chemical series, forms anode and copper automatically becomes cathode. Example: Steel screws in a brass marine hardware, steel pipe connected to copper etc. Underground iron pipelines connected to Zn bar, steel pipe connected to copper plumbing, zinc coating on mild steel. Example of Galvanic Corrosion The mechanisms of differential metal corrosion. 17  Different Types of Electrochemical Corrosion Galvanic Corrosion or Bimetallic Corrosion The primary driving force in galvanic corrosion is a property known as potential difference. When a metal is immersed in an electrolyte, it adopts an electrode potential. The value of the electrode potential for various metals is represented in a table known as the galvanic series. The potential difference between the two metals is, therefore, the difference between their respective electrode potentials as defined in the galvanic series. Requirements for Galvanic Corrosion: In order for galvanic corrosion to occur, three elements are required. 1) Two metals with different corrosion potentials 2) Direct metal-to-metal electrical contact 3) A conductive electrolyte solution (e.g. water) must connect the two metals on a regular basis. The electrolyte solution creates a “conductive path”. This could occur when there is regular immersion, condensation, rain, fog exposure or other sources of moisture that dampen and connect the two metals. If any of these elements is missing, galvanic corrosion cannot occur. If, for example, the direct contact between the two metals is prevented (plastic washer, paint film etc.) or if there is some other interruption in the conductive path, there cannot be galvanic corrosion and each metal will corrode at its normal rate in that service environment. Figure 1 shows examples of conditions that do not meet all requirements for galvanic corrosion. 18  Different Types of Electrochemical Corrosion Differential Aeration Differential Aeration corrosion corrosion (( Concentration Concentration cell cell corrosion) corrosion) It occurs when a metal is exposed to differential air concentration or oxygen concentration. Since the cathodic reaction requires oxygen, the part of the metal exposed to higher concentration of oxygen acts as cathode. On the other hand, the metallic part exposed to lower oxygen concentration acts as anodic region and corrosion occurs at this region. For example, when a zinc rod is partially immersed in neutral salt solution, the metal above the water line is more oxygenated, while the portion that is immersed has smaller oxygen concentration and thus become anodic. Hence a potential difference is created, which causes the flow of current between two differentially aerated areas of same metal. Anode 𝒁𝒏 𝒁𝒏𝟐+ + 𝟐𝒆− Zn-rod Flow of 𝟏 − Cathode 𝑶 + 𝟐𝒆 + 𝑯𝟐 𝑶 𝟐 𝟐 𝟐𝑶𝑯 More oxygenated part (Cathode) Electron NaCl 𝟏 − 𝑶 + 𝟐𝒆 + 𝑯𝟐 𝑶 𝟐𝑶𝑯 Solution 𝟐 𝟐 Corroding The circuit is completed by migration of ions through the electrolyte and − Anode flow of electrons through the metal from anode to cathode. 𝒆 𝒁𝒏𝟐+ 𝒁𝒏 𝒁𝒏𝟐+ + 𝟐𝒆− (Less oxygenated part) 𝒁𝒏𝟐+ ZnCl2 Mechanism of differential aeration corrosion caused by primary immersion of metal 19  Different Types of Electrochemical Corrosion Pitting Corrosion It is defined as intense, localized, accelerated attack resulting in the formation of a pinholes, pits and cavities on the metal surface. Such a type of corrosion takes place when there is a breakdown, peeling or cracking of a protective film due to scratches, abrading action, sliding under load etc. Waterline Corrosion When water is stored in a container or a steel tank, it is generally found that most of the corrosion takes place just beneath the line of water level. The area above waterline is highly oxygenated and acts as cathode, while the area just beneath the waterline is poorly oxygenated and becomes anodic site. This type of corrosion is also a consequence of differential aeration. Ocean-going ships, water storage steel tanks or steel pipes carrying liquid when exposed to atmosphere, undergoes this type of corrosion. The metallic part just below the waterline is more anodic to part above the waterline. Therefore, the metal just below the waterline undergoes corrosion. 20  How Nature of Metals affect the Corrosion? Position in the galvanic series: Metals with low reduction electrode potentials exhibit high reactivity and hence, are more susceptible to get corroded. Thus, the tendency of a metal to get corroded decreases with increase in the reduction potential. Metals, such as Li, Na, Mg, Zn etc. undergo severe corrosion in corrosion environment due to their low reduction potentials. On the other hand, noble metals with high reduction potentials, like Ag, Au, Pt etc. are less susceptible for undergoing corrosion. However, few metals deviate from their normal tendency due to their ability to develop passive layers. Electrode Potential Difference: The potential difference between the anodic and cathodic regions is known as open circuit potential difference (OCPD). Larger is the OCPD, higher is the rate of corrosion. For example, the potential difference between Fe & Ag is 1.14 V and that between Fe & Sn is 0.84 V. Therefore, Fe corrodes faster when it is in contact with Ag than with tin in a corrosion environment. Surface State of the Metal: State or conditions of the surface of the metals also affect the rate of corrosion process. For example, if the surface is covered with sand or dust particles, then severe pitting corrosion occurs. Corrosion at the surface is also accelerated by the presence of water droplets. Heterogeneity at the surface of the metal also results in high corrosion rate. Presence of uneven stress helps in high corrosion. Hydrogen Overvoltage: In most of the electrochemical corrosion process, the competing cathodic reaction involves the liberation of hydrogen gas due to the reduction of hydrogen ions. Hydrogen overvoltage is the measure of the tendency of an electrode to liberate hydrogen gas. A metal with low hydrogen overvoltage is more susceptible to corrosion. Because low hydrogen overvoltage accelerate the cathodic reaction, which in turn increases anodic reaction rate and hence, corrosion rate. Formation of protective film of corrosion products by metals: A few metals show tendency to develop passive, but protective films on exposure to corrosive medium. The passive layers, being stable, highly insoluble with low conductivity, act as a protective layer and prevent further corrosion. Metals like Al, Ti and Cr, develop such a layer on their surface and become passive to corrosion. On the other hand, if the layer of the corrosion products is soluble, non-uniform, volatile or porous, then the corrosion continues to occur. For example, metal oxide layers of Fe & Zn cannot prevent the corrosion process completely as they are porous. Relative area of Anode and Cathode: When the relative area of anode is small compared to that of cathode, corrosion rate becomes high. It is because of the high anodic current density. For example, any gap on the tin coated iron results in severe corrosion due to small anodic area and large cathodic area. 21  How Nature of Corroding Environment affect the Corrosion? pH: Lower is the pH, higher is the corrosion rate. In acidic medium, generally the corrosion rate is more. If the pH < 3, severe corrosion occurs even in the absence of air due to liberation of hydrogen gas at the cathodic area. At pH> 10, corrosion of Fe practically ceases due to formation of protective coating of hydrous oxide. However, metals like Al, Zn undergo high corrosion in highly alkaline solution. Temperature: The rate of any chemical reaction generally increases with increase in temperature and increase in temperature also enhances the conductivity of the medium. As the corrosion process is either chemical or electrochemical reaction, the rate of corrosion increases with increase in the temperature. That means, higher is the temperature, more is the corrosion rate. Conductance of the medium: Electrochemical corrosion involves electron transfer. The presence of conducting medium facilitates the transfer of electrons and hence, increases the corrosion rate. For this reason, the corrosion rate is higher in saline water than in normal water. Humidity: Humidity is the measure of the moisture content in the atmosphere. Rate of corrosion increases with increase in humidity upto certain value, called as critical humidity. But corrosion rate abruptly increases above the critical humidity. 22  How to Prevent Corrosion? Proper designing The design of the material should be such that corrosion, even if it occurs, is uniform and does not result in intense and localized corrosion”. Important design principles are: Avoid the contact of dissimilar metals in the presence of a corroding solution, otherwise the corrosion is localized on the more active metal and less active metal remains protected. 1) When two dissimilar metals are to be in contact, the anodic material should have as large area as possible; whereas the cathodic metal should have as much smaller area as possible. 2) If two dissimilar metals in contact have to be used, they should be as close as possible to each other in the galvanic series. 3) Whenever the direct joining of dissimilar metals is unavoidable, an insulating fitting may be applied in between them to avoid the direct metal to metal contact. 4) The anodic metal should not be painted or coated, when in contact with a dissimilar cathodic metal. 23  How to Prevent Corrosion? 1) A proper design should avoid the presence of crevices between adjacent parts of structure, even in case of the same metal, since crevices permit concentration differences. 2) Sharp corners and recesses should be avoided, as they are favorable for the formation of stagnant areas and accumulation of solids. 3) The equipment should be supported on legs to allow free circulation of air and prevent the formation of stagnant pools or damp areas. 24  How to Prevent Corrosion? Use of pure metal Impurities in a metal cause heterogeneity, which decrease corrosion resistance of the metal. Hence corrosion resistance of any metal is improved by increasing its purity. e.g. Al, Mg. ; the corrosion resistance of Al depends on its oxide film formation, which is highly protective only on the high purity metal. Using metal alloys Corrosion resistance of most metals is best increased by alloying them with suitable elements. For maximum corrosion resistance, the alloy should be completely homogeneous. 25  How to Prevent Corrosion? Cathodic protection The principle involved here is to force the metal to be protected as to behave like a cathode. There are two types of cathodic protections. I. Sacrificial anodic protection method The metallic structure to be protected is connected by a wire to the more anodic metal, so that active metal itself get corroded slowly, while the parent structure is protected. The more active metal is called “sacrificial anode”, which must be replaced, when consumed completely. Metals commonly used as sacrificial anodes are Mg & Zn. 26  How to Prevent Corrosion? II. Impressed current cathodic protection: An impressed current is applied in opposite direction to nullify the corrosion current, and convert the corroding metal from anode to cathode. Usually a sufficient D.C. is applied to an insoluble anode, buried in the soil and connected to the metallic structure to be protected. The anode is usually in a backfill (composed of cock breeze or gypsum), so as increase the electrical contact with the surrounding soil. This kind of protection technique is useful for large structures for long term operations. 27  How to Prevent Corrosion? Use of inhibitors A corrosion inhibitor is “a substance when added in small quantities to the aqueous corrosive environment, effectively decreases the corrosion of the metal”. i) Anodic inhibitors: Anodic inhibitors stop the corrosion reaction, occurring at anode, by forming a precipitate with a newly produced metal ion. These are adsorbed on the metal surface in the form of a protective film or barrier. Examples are chromates, phosphates, tungstates and other transition metals with high oxygen content. ii) Cathodic inhibitors: In acidic solutions, the main cathodic reaction is evolution of hydrogen. 2𝐻 + 𝑎𝑞 + 2𝑒 − → 𝐻2 Corrosion may be reduced either by slowing down the diffusion of hydrated H+ ions to the cathode and/or by increasing the over voltage of hydrogen evolution. The diffusion of H+ ions is considerably decreased by organic inhibitors like amines, mercaptans, heterocyclic nitrogen compounds, substituted urea and thiourea, heavy metal soaps, which are capable of being adsorbed at metal surfaces. b) In neutral solutions, the cathodic reaction is 𝐻2 𝑂 + 𝑂2 + 2𝑒 − = 𝑂𝐻− (𝑎𝑞) Corrosion is controlled either by eliminating oxygen from the corroding medium or by retarding its diffusion to the cathodic areas. The oxygen is eliminated either by reducing agents (like Na2SO3) or by de-aeration. The inhibitors like Mg, Zn or Ni salts tend to retard the diffusion of 𝑂𝐻 − ions to cathodic areas. 28  How to Prevent Corrosion? Protective coatings It is the oldest of the common procedures for corrosion prevention. A coated surface isolates the underlying metal from the corroding environment. i) The coating applied must be chemically inert to the environment under particular conditions of temperature and pressure. ii) The coatings must prevent the penetration of the environment to the material, which they protect. There are mainly three types of protective coatings a) Metallic coatings: b) Inorganic coatings (chemical conversion) c) Organic coatings (paints etc.,) 29  How to Prevent Corrosion? Metallic coatings: A metal is coated on the other metal, in order to prevent corrosion. These are of two types I) Anodic coatings: These are produced from coating-metals, which are “anodic” to the base metal. This provides the complete protection to the underlying base metal as long as the coating intact. However, the formation of the pores or cracks on the protective layer can set up severe galvanic corrosion leading to complete destruction of the base metal. E.g.: In case of galvanized steel, zinc, the coating-metal being anodic is attacked; leaving the underlying cathodic metal (iron) unattacked (Figure 17 ) 30  How to Prevent Corrosion? Cathodic coatings These are obtained by coating a more noble metal having higher electrode potential than the base metal. The cathodic coating provides effective protection to the base metal only when they are completely continuous and free from pores, breaks or discontinuities. An example of cathodic coating is Tinning, coating of tin on iron (Figure 18 ). 31