CHAPTER I 1.0 INTRODUCTION

Transcription

1 CHAPTER I INTRODUCTION 1.1. The Nature of Matter [1] All matter is made of atoms composed of protons, neutrons, and electrons. The center, or nucleus, of the atom is composed of positively charge protons and neutral neutrons. The outside of the atom has negatively charged electrons in various orbits. All atoms have the same number of protons (positively charged) and electrons (negatively charged). Therefore all atoms have a neutral charge (the positive and negative charges cancel each other). Most atoms have approximately the same number of neutrons as they do protons or electrons, although this is not necessary, and the number of neutrons does not affect the identity of the element. The number of protons (atomic number) in an atom determines which kind of atom we have, and the atomic mass (weight) of the atom is determined by the number of protons and neutrons in the nucleus (the electrons are so small as to be almost weightless). There are over 100 different elements that have been discovered. The letter symbols for the elements come from their Latin names, so for example, H stands for hydrogen, C for Carbon, O for oxygen, while Fe stands for iron and Cu stands for copper. The Periodic Table of the Elements is shown in Fig.1.1.

2 4 Fig Periodic Table of the Elements. 1.1 (a) Ions: Ions are formed when atoms, or groups of atoms, lose or gain electrons.metals lose some of their electrons to form positively charged ions, e.g.fe +2, Al +3, Cu +2, etc. Nonmetals gain electrons and form negatively charged ions, e.g.cl -, O -2, S -2, etc. 1.1 (b) Molecules: Compounds are groups of metals and nonmetals that form distinct chemicals. Most of us are familiar with the formula H 2 O, which indicates that each water molecule is made of two hydrogen atoms and one oxygen atom. Many molecules are formed by sharing electrons between adjacent atoms. A water molecule has adjacent hydrogen and oxygen atoms sharing some of their electrons. Only a few of the elements are common, and most corrosion occurs due to metallic elements (iron, aluminum, copper, zinc, etc.) reacting with common nonmetallic elements (oxygen, chlorine, sulfur, etc.).

3 1.2. Historical background of corrosion 5 Corrosion can be viewed as a Universal Phenomenon, Omnipresent and Omnipotent. It is there every where, air, water, soil and in every environment, we encounter. The word corrosion is as old as the earth, but it has been known by different names. Corrosion is known commonly as rust, an undesirable phenomena which destroys the luster and beauty of objects and shortens their life. A Roman philosopher, Pliny (AD 23 79) wrote about the destruction of iron in his essay Ferrum Corrumpitar. Corrosion since ancient times has affected not only the quality of daily lives of people, but also their technical progress. There is a historical record of observation of corrosion by several writers, philosophers and scientists, but there was little curiosity regarding the causes and mechanism of corrosion until Robert Boyle wrote his Mechanical Origin of Corrosiveness. Philosophers, writers and scientists observed corrosion and mentioned it in their writings: Pliny the elder (AD 23 79) wrote about spoiled iron. Herodotus (fifth century BC) suggested the use of tin for protection of iron. Lomonosov ( ). Austin (1788) noticed that neutral water becomes alkaline when it acts on iron. Thenard (1819) suggested that corrosion is an electrochemical phenomenon. Hall (1829) established that iron does not rust in the absence of oxygen. Davy (1824) proposed a method for sacrificial protection of iron by zinc. De la Rive (1830) suggested the existence of microcells on the surface of zinc. The most important contributions were later made by Faraday ( ) [2] who established a quantitative relationship between chemical action and electric current. Faraday s first and second laws are the basis for calculation of corrosion rates of metals. Ideas on corrosion control started to be generated at the beginning of nineteenth century. Whitney (1903) provided a scientific basis for corrosion control based on electrochemical observation. As early as in eighteenth century it was observed that iron corrodes rapidly in dilute nitric acid but remains unattacked in concentrated nitric acid. Schönbein in 1836 showed that iron could be made passive [3]. It was left to U. R. Evans to provide a modern understanding of the causes and control of corrosion based on his classical electrochemical theory in Considerable progress

4 6 towards the modern understanding of corrosion was made by the contributions of Evans [4], Uhlig [5] and Fontana [6]. The above pioneers of modern corrosion have been identified with their well known books in the references given at the end of the chapter. Corrosion laboratories established in M.I.T, USA and University of Cambridge, UK, contributed significantly to the growth and development of corrosion science and technology as a multi disciplinary subject. In recent years, corrosion science and engineering has become an integral part of engineering education globally Definitions In nature, metals are not found in free state due to their reactivity. It is the ore from which the metals are extracted by metallurgical process. Metallurgy requires a large amount of heat energy[7]. The metals are thermodynamically unstable in their free state. They are stable in the form of certain compounds. This is given as: Metallurgy Corrosion Ore Metal Corrosion product (thermodynamically (thermodynamically (thermodynamically Stable) unstable) stable) Corrosion is a process of formation of the compound of pure metal by the chemical reaction between metallic surface and its environment. It is an oxidation process. It causes loss of metal. Hence, disintegration of a metal by its surrounding chemicals through a chemical reaction on the surface of the metal is called corrosion. Example: Formation of rust on the surface of iron, formation of green film on the surface of copper.the responsible factors for the corrosion of a metal are the metal itself, the environmental chemicals, temperature and the design. Corrosion is a natural and costly process of destruction like earthquakes, tornados, floods and volcanic eruptions, with one major difference. Whereas we can be only a silent spectator to

5 7 the above processes of destruction, corrosion can be prevented or at least controlled. Several definitions of corrosion have been given and some of them are reproduced below: environment. Corrosion can be defined as the degradation of a material due to a reaction with its Degradation implies deterioration of physical properties of the material. This can be a weakening of the material due to a loss of cross-sectional area, it can be the shattering of a metal due to hydrogen embrittlement, or it can be the cracking of a polymer due to sunlight exposure.most corrosion of metals is electrochemical in nature. Corrosion of metals is defined as the spontaneous destruction of metals in the course of their chemical, electrochemical and biochemical interactions with the environment [6,8]. Thus it is exactly the reverse of extraction of metal from ores. In most of the environments the metals are not stable, but tend to revert to compounds which are more stable, a process which is called corrosion. Corrosion is derived from the Latin word corrosus, meaning, gnawed away. The same amount of energy needed to extract metals from their minerals is emitted during the chemical reactions that produce corrosion [9]. Useful chemical properties are lost by chemical or electrochemical reaction of a metal with its environments. The best known examples of metallic corrosion are rusting of iron and steel, tarnishing of silver and copper, dulling of brass and fogging of nickel. Uhlig and Wistong [5] describe corrosion as the reaction of a solid with its environment It is a partial degeneration from the metal stable condition of the metal to stable condition of the mineral, accompanied by the decrease in the free energy of the system [10]. Material other than ceramics, plastics and concrete may also corrode. Corrosion returns the metal to its combined state in chemical compounds that are similar or even identical to the minerals from which the metals were extracted. Thus corrosion may also be called as extractive metallurgy in the reverse. In the present days it is necessary to pay more attention to corrosion because of the increasing use of metals in all fields of technology, the increase in air and water pollution resulting in a more corrosive environment. Despite different definitions, it can be observed that corrosion is basically the result of interaction between materials and their environment. Up to the 1960s, the term corrosion was

6 8 restricted only to metals and their alloys and it did not incorporate ceramics, polymers, composites and semiconductors in its regime. The term corrosion now encompasses all types of natural and man-made materials including biomaterials and nanomaterials, and it is not confined to metals and alloys alone. The scope of corrosion is consistent with the revolutionary changes in materials development witnessed in recent years Cost of corrosion Corrosion is recognized as one of the most serious problems in out modern society. Cost of corrosion studies have been undertaken by several countries including the United States, the United Kingdom, Australis, Japan, Finland, Kuwait, Germany, Sweeden, India and China. These studies have ranged from and extensive efforts to informal and modest efforts. The economic cost of corrosion world wide is enormous and runs into billions of dollars. It was estimated that one ton of steel runs into rust every 90 seconds [11], and considerable portion of steel produced are used for replacement of corroded parts. Frequently we come across expense Fig directly relating to the economic cost of corrosion. But indirect costs resulting from actual possibilities of corrosion are more difficult to evaluate and are probably ever greater. According to the National Association of Corrosion Engineers international India [12] section (NACE), the annual direct cost of corrosion in India may be 4% of GNP (Gross National Product), which is estimated to be Rs.350 billion per year The solution to reduce this phenomenal loss is to ensure that industries take up corrosion prevention as an important issue even at the design stage. The losses due to corrosion may be divided into direct and indirect [13], as detailed below. Some of the indirect costs are 1) plant down time 2) loss of product 3) loss of efficiency 4) contamination and 5) over design. It is apparent that a person working in the field of corrosion is responsible not only for the protection of the product, equipment and welfare of the individuals but also for providing this at a reasonable cost The flow chart of Concentration cell corrosion is shown in Fig. 1.2.

7 9 Cost of corrosion Direct losses Indirect losses 1) Over design 1. Economical 2) Inability to use other a) Contamination of the desirable materials product (b) Loss 3) Cost of repair or replacement of valuable Of corroded equipment production c) Loss of productions Fig Cost of corrosion 2. Social Safety Eg.Sudden failure of equipment, fire explosion or release of toxic products Corrosive Environment [14-30] Corrosion cannot be defined without a reference to environment. All environments are corrosive to some degree. Following is the list of typical corrosive environments: Air and humidity. Fresh, distilled, salt and marine water. Natural, urban, marine and industrial atmospheres. Steam and gases, like chlorine. Ammonia. Hydrogen sulfide. Sulfur dioxide and oxides of nitrogen. Fuel gases.

8 10 Acids. Alkalies. Soils. It may, therefore, be observed that corrosion is a potent force which destroys economy, depletes resources and causes costly and untimely failures of plants, equipment and components Factors influencing corrosion [7] Since corrosion is a process of destruction of metal surface by its environment, the two factors that govern the corrosion process are: (i) (ii) Metallic and Environmental. (i) Nature of metal: Different properties of a metal are responsible for corrosion. These properties are given here. (a) Position of metal in galvanic series : It decides the corrosion rate. A metal having higher position in galvanic series undergoes corrosion when con- nected to another metal below it. Also, more difference in the position of galvanic series will cause faster corrosion at anodic metal. (b) Hydrogen over voltage : In case of zinc metal placed in a normal solution of H 2 SO 4, reaction takes place forming bubbles of hydrogen gas on zinc surface. The process is slow due to high hydrogen over voltage of zinc (0.76 V). The addition of few drops of CuSO 4 accelerates corrosion due to reduction of hydrogen over voltage (0.34V). Further, faster corrosion is observed in the presence of PtCl 4 (hydrogen over voltage = 0.2 V). The reduction in over voltage of corroding metal or alloy accelerates the rate of corrosion. Hence hydrogen over voltage governs the process of corrosion. (c) Purity of metal: Pure metal resists corrosion, while impurities in a metal form a local galvanic cell (metal as anode and impurity as cathode) and result in the corrosion of metal. Rate of corrosion increases due to more exposure of impurities. For alloys the system is a homogeneous solid solution, hence no local action and no corrosion.

9 (d) (e) (f) 11 Relative areas of anode and cathode: Smaller the area of anode com- pared to cathode will lead to faster corrosion of anode. It is because the corrosion current at anode and at cathode will be same. But for small anodic area the current density will be large at anode and larger cathodic area will demand more electron which will be fulfilled by fast reaction at anode (oxdidation), i.e. rapid corrosion. Physical state of the metal: Small granular metal will corrode faster than the larger one. Also the type of structure formed by a metal will have effect on the corrosion rate. A bent metal (stress) is rapidly corroded due to stress. Nature of oxide film: An oxide film is formed by the reaction between metal and oxygen. If this oxide film is porous and oxygen can be diffused through it, more corrosion is expected (already shown in dry or chemical corrosion). Also, if volume of metal oxide is more than the volume of metal (The specific volume ratios of Ni, Cl, W are 1.6, 2.0 and 3.6) least corro- sion or no further corrosion occurs. (g) Volatility and solubility of corrosion product: In both the cases, the corrosion will be faster. MnO3, SnCl4 are volatile, so faster is corrosion of Sn in chlorine atmosphere. In case of soluble corrosion product, it will be enhanced by water and metal surface will be exposed for further corrosion. The role of environment in the corrosion of a metal is very important. Environmental parameters like temperature, humidity, ph, etc. play important role. The effect is discussed here. (a) (b) (c) Temperature: The rate of diffusion increases by rise in temperature, hence the rate of corrosion is also increased. At higher temperature, passive metals also become active and undergo corrosion. But higher temperature reduces the concentration of oxygen and hence corrosion is reduced (in case of water where oxygen is dissolved). Humidity: In humidity, gases like CO 2, SO 2, NO x are dissolved which form electrolytes. It will cause galvanic corrosion. Some oxides are water soluble, humidity washes away the corrosion products and metal surface is further corroded. Other soluble corrosion products can also be washed away by humidity, causing further corrosion. Impurity of atmosphere: Pollutants like H 2 S, SO 2, CO 2 and acid vapours cause more pollution where they dissolve. In sea water (salty in nature which acts as an electrolyte)

10 12 corrosion rate increases. Some suspended particles are dissolved in humidity and form electrolyte which helps in corrosion. (d) (e) (f) (g) ph value: ph value means concentration of H+ (acidic nature). In acidic medium (ph less than 7), corrosion is faster. Also, in basic medium ph > 7, some metals such as Pb, Zn, Al, etc. form complexes and hence they corrode. Pourbiax relation between ph of medium and potential of metal deals with the corrosion process and it gives idea how to reduce corrosion. Example: Zn corrodes minimum at ph 11, but at higher ph (more than 11) it corrodes faster. At ph 5.5, Al corrodes minimum. Nature of ions present: Cu ++ ions present in the vicinity of Fe, accelerate corrosion, while silicates present in the vicinity resist corrosion. Conductance effect: Due to presence of salts and water in earth, it is of con- ducting nature. More conductance leads to more stray current and hence fast corrosion. Dry sandy soil is less conducting and hence less corrosion, while mineralised clay soil is more conducting hence more corrosion occurs. Oxygen concentration and oxygen concentration cell: Oxygen is one of the important element responsible for corrosion. It forms oxides and hydroxides (in presence of H 2 O) on the surface of metal as corrosion product. Oxygen concentration cell is formed on the surface of metal due to difference in oxygen concentration (iron rod half dipped in water corrodes due to this effect). Dipped portion will be anode and outer por- tion will be cathode Inter-disciplinary Nature of Corrosion [14-30] The subject of corrosion is inter-disciplinary and it involves all basic sciences, such as physics, chemistry, biology and all disciplines of engineering, such as civil, mechanical, electrical and metallurgical engineering Reasons to study corrosion [14-30] Corrosion studies have become important due to increasing awareness of the need to conserve the world s metal resources. The reasons for the important or studies on corrosion are:

11 13 Several engineering disasters, such as crashing of civil and military aircraft, naval and passenger ships, explosion of oil pipelines and oil storage tanks, collapse of bridges and decks and failure of drilling platforms and tanker trucks have been witnessed in recent years. Corrosion has been a very important factor in these disasters. Applying the knowledge of corrosion protection can minimize such disasters. The designing of artificial implants for the human body requires a complete understanding of the corrosion science and engineering. Surgical implants must be very corrosion-resistant because of corrosive nature of human blood. Materials are precious resources of a country. Our material resources of iron, aluminum, copper, chromium, manganese, titanium, etc. are dwindling fast. Some day there will be an acute shortage of these materials. An impending metal crisis does not seem anywhere to be a remote possibility but a reality. There is bound to be a metal crisis and we are getting the signals. To preserve these valuable resources, we need to understand how these resources are destroyed by corrosion and how they must be preserved by applying corrosion protection technology. Corrosion is a threat to the environment. For instance, water can become contaminated by corrosion products and unsuitable for consumption. Corrosion prevention is integral to stop contamination of air, water and soil. Engineering knowledge is incomplete without an understanding of corrosion. Aeroplanes, ships, automobiles and other transport carriers cannot be designed without any recourse to the corrosion behaviour of materials used in these structures Different forms of corrosion There are several forms of corrosion, also referred to as modes or mechanisms of corrosion, Sub-forms can also be identified for several forms of corrosion [6,31]. The different forms and sub-forms of corrosion are analogous to different failure mechnisams in the mechanical world (fast fracture, fatigue, ductile tearing, brittle cleavage, wear etc..,)

12 Uniform Corrosion[32-40] :- This is also called general corrosion. The surface effect produced by most direct chemical attacks (e.g., as by an acid) is a uniform etching of the metal. The picture of Uniform Corrosion is shown in Fig Fig 1.3. Uniform Corrosion Galvanic Corrosion[32-40] :- Galvanic corrosion is an electrochemical action of two dissimilar metals in the presence of an electrolyte and an electron conductive path. It occurs when dissimilar metals are in contact. The picture of Galvanic Corrosionis shown in Fig Fig Galvanic Corrosion Concentration Cell Corrosion[32-40]:- Concentration cell corrosion occurs when two or more areas of a metal surface are in contact with different concentrations of the same solution. The picture of Concentration cell corrosion is shown in Fig Fig Concentration Cell Corrosion

13 Pitting Corrosion [32-40] :- Pitting corrosion is localized corrosion that occurs at microscopic defects on a metal surface. The pits are often found underneath surface deposits caused by corrosion product accumulation. The picture of Pitting Corrosion is shown in Fig Fig 1.6. Pitting Corrosion Crevice Corrosion [32-40] Crevice or contact corrosion is the corrosion produced at the region of contact of metals with metals or metals with nonmetals. It may occur at washers, under barnacles, at sand grains, under applied protective films, and at pockets formed by threaded joints. The picture of Crevice Corrosionis shown in Fig Fig 1.7.Crevice Corrosion

14 Filiform Corrosion[32-40] This type of corrosion occurs on painted or plated surfaces when moisture permeates the coating. Long branching filaments of corrosion product extend out from the original corrosion pit and cause degradation of the protective coating. The picture of Filiform Corrosion is shown in Fig Fig 1.8.Filiform Corrosion Intergranular Corrosion[32-40] Intergranular corrosion is an attack on or adjacent to the grain boundaries of a metal or alloy. The picture of 7 Intergranular Corrosionis shown in Fig Fig 1.9.Intergranular Corrosion

15 Stress Corrosion Cracking[32-40] Stress corrosion cracking (SCC) is caused by the simultaneous effects of tensile stress and a specific corrosive environment. Stresses may be due to applied loads, residual stresses from the manufacturing process, or a combination of both. The picture of Stress Corrosion Crackingis shown in Fig Fig 1.10.Stress Corrosion Cracking Corrosion Fatigue[32-40] Corrosion fatigue is a special case of stress corrosion caused by the combined effects of cyclic stress and corrosion. No metal is immune from some reduction of its resistance to cyclic stressing if the metal is in a corrosive environment. The picture of Corrosion Fatigueis shown in Fig Fig 1.11.Corrosion Fatigue

16 Fretting Corrosion[32-40] The rapid corrosion that occurs at the interface between contacting, highly loaded metal surfaces when subjected to slight vibratory motions is known as fretting corrosion. The picture of 10 Fretting Corrosionis shown in Fig Fig 1.12.Fretting Corrosion Erosion Corrosion[32-40] Erosion corrosion is the result of a combination of an aggressive chemical environment and high fluid-surface velocities. The picture of Erosion Corrosionis shown in Fig Fig 1.13.Erosion Corrosion

17 Dealloying[32-40] Dealloying is a rare form of corrosion found in copper alloys, gray cast iron, and some other alloys. Dealloying occurs when the alloy loses the active component of the metal and retains the more corrosion resistant component in a porous "sponge" on the metal surface. The picture of Dealloyingis shown in Fig Fig 1.14.Dealloying Hydrogen Damage[32-40] Hydrogen embrittlement is a problem with high-strength steels, titanium, and some other metals. Control is by eliminating hydrogen from the environment or by the use of resistant alloys. The picture of Hydrogen Damageis shown in Fig Fig 1.15.Hydrogen Damage

18 Corrosion in Concrete [32-40] Concrete is a widely-used structural material that is frequently reinforced with carbon steel reinforcing rods, post-tensioning cable or prestressing wires. The steel is necessary to maintain the strength of the structure, but it is subject to corrosion. The picture of Corrosion in Concrete is shown in Fig Fig 1.16.Corrosion in Concrete Microbial Corrosion [32-40] Microbial corrosion (also called microbiologically -influenced corrosion or MIC) is corrosion that is caused by the presence and activities of microbes. This corrosion can take many forms and can be controlled by biocides or by conventional corrosion control methods. The picture of Microbial Corrosionis shown in Fig Fig Microbial Corrosion

19 1.10 Classfiction of Corrosion 21 Depending upon the nature of corrosion and the factors affecting it, corrosion is classifieds as : Dry or chemical corrosion In dry corrosion, the metal is surrounded by gases such as oxygen, halogens, sulphur dioxide, hydrogen sulphide, nitrogen etc., in the surrounding environment and as a result, corrosion ocfurs mainly through the direct chemical action of environmental or atmospheric gases with metal surfaces in immediate proximity. This type of corrosion produces two important effects on the metal. Metal is consumed. The properties of the metal are changed The extent of dry corrosion depends upon the following factors Chemical affimity between the corrosive environment and solid metals Capacity or ability of reaction product on metal surface to form a protective film Wet or electrochemical corrosion This type of corrosion occurs when a conducting liquid is in contact with metals or when two dissimilar metals or alloys are either immersed or dipped partially in a solution. At anodic area, oxidation reaction (i.e). liberation of free electrons takes place, so, anodic metal is destroyed. At anode : M M n+ +ne - (oxidation) 1.1. At cathodic area. reduction reaction (i.e). gain of electorns takes place. Usually cathode rections do not affect the cathode metal. At catode : nh + +ne - nh (reduction) 1.2. The electrochemical corrosion involves The formation of anodic and cathodic area Presence of a conducting medium Corrosion of anodic areas.

20 22 Formation of corrosion product Principles of corrosion Corrosion resistance or chemical resistance depends on many factors It s completed and comprehensive study requires an understanding of several field of scientific knowledge as indicated in Fig Thermodynamics are electrochemistry are of great importance in understanding and controlling corrosion [6.13]. Metallurgical principles Thermodynamic principles Corrosion principles Electrochemical principles Physical and chemical principles Fig Principles of corrosion Different theories of corrosion [7] There are three theories of corrosion: Acid theory Dry or chemical corrosion Galvanic or electrochemical or wet corrossion Acid theory of corrosion This theory suggests that corrosion of a metal (iron) is due to the presence of acids surrounding it. According to this theory, iron is corroded by atmospheric carbon di- oxide, moisture and

21 23 oxygen. The corrosion products are the mixture of Fe(HCO3)2, Fe(OH)CO3 and Fe(OH)3. The chemical reactions suggested are given below 1 Fe + 2Co 2 + H 2 O + O 2 Fe(HCO 3 ) Fe(HCO 3 ) 2 + H 2 O + O 2 2Fe(OH)CO 3 + 2CO H 2 O Fe(OH)CO 3 + 2H 2 O 2Fe(OH) 3 + 2CO This theory is supported by the analysis of rust that gives the test for CO= ion. Further, the process of rusting is reduced by the presence of lime and caustic soda (these two can absorb CO 2, thus reducing corrosion) Chemical theory of corrosion According to this theory, corrosion on the surface of a metal is due to direct reaction of atmospheric gases like oxygen, halogens, oxides of sulphur, and oxides of nitrogen, hydrogen sulphide and fumes of chemicals with metal. The extent of corrosion of a particular metal depends on the chemical affinity of the metal towards reactive gas. Oxygen is mainly responsible for the corrosion of most metallic substances when compared to other gases and chemicals. There are three main types of dry corrosion. Oxidation corrosion (Reaction with oxygen): Some of the metals directly react with oxygen in the absence of moisture. Alkali and alkaline earth met- als react with oxygen at room temperature and form corresponding oxides, while some metals react with oxygen at higher temperature. Metals like Ag, Au and Pt are not oxidised as they are noble metals.

22 24 M M n+ + ne - (if n = 1) O 2 + 2e - O - for alkali metals M + O 2 M 2 O Oxidation of a metal, metal oxide is formed as a thin film on the metallic surface is given in Fig Fig Oxygen from Atmosphere During oxidation of a metal, metal oxide is formed as a thin film on the metallic surface which protects the metal from further corrosion. If diffusion of either oxygen or metal is across this layer, further corrosion is possible. Thus, the layer of metal oxide plays an important role in the process of corrosion. Oxides of Pb, Al and Sn are stable and hence inhibit further corrosion. They form a stable, tightly adhering oxide film. In case of porous oxide film, atmospheric gases pass through the pores and react with the metal and the process of corrosion continues to occur till the entire metal is converted into oxide. Porous oxide layer is formed by al- kali and alkaline earth metals. Molybdenum forms a volatile oxide film of MoO 3 which accelerates corrosion. Au, Ag, Pt form unstable oxide layer which decomposes soon after the formation, thereby preventing further corrosion. Pilling Bedworth Rule: If volume of metal oxide on the surface of a metal is more than or equal to the volume of metal, the oxide layer will be protective. For example Al 2 O 3, Fe, Ni, ZnW, Cr will be non-protective if volume of oxide is less than volume of metal. (The specific

23 25 volume ratio of W is 3.6, Cr = 2.0, Ni = 1.6. Hence, the rate of corrosion is very less in tungsten.) It is called Pilling Bedworth rule. Corrosion by other gases such as Cl 2, SO 2, H 2 S, and NOx: In dry atmosphere, these gases react with metal and form corrosion products which may be protective or non-protective. Dry Cl 2 reacts with Ag and forms AgCl which is a protective layer, while SnCl 4 is volatile. In petroleum industries at high temperatures, H 2 S attacks steel forming FeS scale which is porous and interferes with normal operations. Liquid metal corrosion: In several industries, molten metal passes through metal- lic pipes and causes corrosion due to dissolution or due to internal penetration. For example, liquid metal mercury dissolves most metals by forming amalgams, thereby corroding them Wet or electrochemical theory of corrosion by taking rusting of iron as example It is a common type of corrosion of metal in aqueous corrosive environment. This type of corrosion occurs when the metal comes in contact with a conducting liquid or when two dissimilar metals are immersed or dipped partly in a solution. According to this theory, there is the formation of a galvanic cell on the surface of metals. Some parts of the metal surface act as anode and rest act as cathode. The chemical in the environment and humidity acts as an electrolyte. Oxidation of anodic part takes place and it results in corrosion at anode, while reduction takes place at cathode. The corrosion product is formed on the surface of the metal between anode and cathode. To understand the wet theory, let us take the example of corrosion of iron. Oxida- tion of metal takes place at anode while the reduction process takes place at cathode. By taking rusting of iron as an example, the reaction can be explained as that it may occur in two ways: (i) evolution of hydrogen and (ii) (ii) absorption of oxygen. At anode: oxidation occurs. Fe Fe ++ +2e - (oxidation) 1.9.

24 26 At cathode: Case I: Evolution of H 2 The hydrogen ions (H + ) are formed due to the acidic environment and the following reaction occurs in the absence of oxygen 2H+ + 2e H 2 (reduction) The overall reaction is Fe + 2H + Fe +2 + H 2 In this case, metals react in the acidic environment and are dissolved (undergo corrosion) to release H 2 gas. All metals above hydrogen in electrochemical series can show this type of corrosion. In hydrogen evolution type of corrosion, anodic area is large as compared to its cathodic area. Mechanism of wet corrosion by hydrogen evolution is given in Fig. 1.20(a), Fig (a) Mechanism of wet corrosion by hydrogen evolution

25 27 Mechanism of wet corrosion by oxygen absorption is given in Fig. 1.20(b). Fig (b) Mechanism of wet corrosion by oxygen absorption Case II: Absorption of O 2 This type of corrosion takes place in neutral or basic medium in the presence of oxygen. The oxide of iron covers the surface of the iron. The small scratch on the surface creates small anodic area and rest of the surface acts as cathodic area. The following chemical reactions occur at anode and cathode. At anode Fe Fe e (oxidation) At cathode O 2 + H 2 O + 2e - 2OH - (reduction) 1.11 Fe+O 2 +H 2 O Fe OH - or Fe (OH) Ferric hydroxide is actually hydrated ferric oxide, Fe 2 O 3.H 2 O, which is a yellowish rust. Anhydrous magnetite, Fe 3 O 4 [a mixture of (FeO + Fe 2 O 3 )], is also formed, which is brown-black in colour. It is markable that the corrosion occurs at anode but the corrosion product is formed near cathode. It is because of the rapid diffusion of Fe ++ as compared to OH. Hence corrosion occurs at anode, but rust is deposited at or near cathode.

26 28 Table 1.1. The differences between dry and wet corrosion Dry corrosion Corrosion occurs in the absence of moisture. It involves direct attack of chemicals on the metal surface. The process is slow. Corrosion products are produced at the site of corrosion. The process of corrosion is uniform. Wet or electrochemical corrosion Corrosion occurs in presence of conducting medium. It involves formation of electrochemical cells. It is a rapid process. Corrosion occurs at anode but rust is deposited at cathode. It depends on the size of the anodic part of metal Galvanic series [7] Electrochemical reactions are predicted by electrochemical series. A metal having higher position can replace (reduce) other metals that have lower position in the series. For example, Zn + CuSO 4 ZnSO 4 + Cu 1.13 that is, Zn + Cu ++ Zn ++ + Cu 1.14 or in other words, zinc will corrode faster than copper. Some exceptions have been observed in this generalisation. For example, Ti is less reactive than Ag. Galvanic series is the series of metals that is made keeping in view the process of corrosion of a metal in a particular atmosphere, i.e. sea water. In galvanic series, oxidation potential of metals is arranged in the decreasing order of activity of a series of metals. The series is towards the increasing noble nature. More anodic: Mg, Mg alloys, Zn, Al, Cd, Fe, Pb, Sn, Ni Mo Fe alloys) Brasses, Cu, Ni, Cr steel alloy, Ag, Ti, Au, Pt towards noble nature.

27 29 Table 1.2. Comparison between Galvanic Series Vs Electrochemical Series Galvanic Series 1. It predicts the corrosive tendencies of metal alloys 2. Calomel electrode is used as a reference electrode 3. Positioning of metal or alloy may change 4. The metals and alloys are immersed in the sea water for study 5. Electrode potentials are measured for both metals and alloys. Electrochemical Series It predicts the relative displacement tendencies Standard hydrogen electrode is used as reference electrode Position of metal is fixed. That cannot be changed concentration of salts of the same metal that was being used Electrode potentials measured only for metals and nonmetals Protection from corrosion [7] Due to corrosion, there is a great loss of material and money. Therefore, it is essential to protect metals from corrosion. Since, there are two components in- volved in corrosion the metal and environment both are considered in corro- sion protection. Following methods have been adopted for the protection of metal from corrosion. Proper designing of an object helps in prevention of corrosion. Contact of two dissimilar metals must be avoided. If it is unavoidable, an- odic area should be very large compared to cathodic area. Two different metals used in the structure should be such that they are oc- cupying near positions in galvanic series. Putting an insulator between two metals resists corrosion. As far as possible, metal used in a structure should be extremely pure. Small amount of impurity causes corrosion. While using an alloy, it should be completely homogeneous. Design or fabricate equipment or metal parts in such a manner that they have minimised sharp edges and corners and also avoid, as for as possible, the crevices in joints, etc. The modification of environment also helps in protection from corrosion. It includes: De-aeration removes oxygen by adjusting temperature and mechanical ageing.

28 30 Deactivation involves addition of chemicals such as Na 2 SO 3, NH 2 NH 2 which absorb oxygen. Dehumidification of environment is done by adding alumina or silica gel. These chemicals absorb humidity from metallic surface. In humidity, gases such as CO 2, H 2 S, SO 2 and HCl give acidic medium responsible for corrosion. They are neutralised by NH 3 or NaOH or lime. There are some chemicals which reduce the rate of corrosion. These chemicals are called corrosion inhibitors. They are of two types, anodic and cathodic. Phosphate, chromate, tungstate protect anode. They form sparingly soluble products which are adsorbed on the surface of metal and hence check corrosion. Cathodic protection is done by organic amines, mercaptans, thiourea and substituted urea. The above chemicals retard reduction reaction taking place at cathode. Also, by the use of salts of Mg, Zn or Ni, the insoluble hydroxides of Mg, Zn, Ni are deposited preventing corrosion. 1 At cathode : H 2 O + O 2 + 2e - 2OH Mg OH - Mg (OH) (precipitate) Methods of application of metal coatings: [7] Metallic dipping causes a coat of metal over the base metal. It is of two types, anodic coating and cathodic coating (a)anodic coating is done by a metal which is situated prior to the base metal in electrochemical series. Coating of zinc over iron is anodic coating and this process has its special name galvanisation. Zinc occupies a position before iron. In this process base metal is immersed in molten coating metal. The base metal should by very clear. Anodic coating of iron is done by Al and Cd also. If any crack or pit is formed on the coated metal, Zn will act as anode and Fe as cathode and hence Fe will not corrode. Since Zn is toxic, galvanised utensils are not used for cooking or food storing. Zine coating protects iron sacrificially. Galvanised iron is used for buckets, tubes, wires, roof sheets etc. The picture of Anodic Coating is given Fig

29 31 Fig Anodic coating In galvanised steel, Zn serves as anode; while iron of steel serves as the cathode. Therefore, the iron is protected, even if it is exposed, when a part of the zinc coating is scrapped off (b)Cathodic coating of a base metal (say iron) is done by another metal which occupies a position below it in electrochemical series but it will be higher corrosion resistant. Coating of Sn over iron (tinning) is an example of cathodic protection. Cathodic protection of iron is successful only when there is no pit or crack formation on the surface of metal. Coating of tin is called tinning. Tin-coated iron is used for cooking as well as stor- age purposes. The picture of Cathodic Coating is given Fig Fig Cathodic coating Tin plated steel. Tin protects the iron, when the coating is continuous. When the coating is broken, the iron of the steel becomes the anode and is subjected to accelerated local corrosion.

30 32 Table 1.3. Differences between Galvanizing &Tinning Galvanizing 1. Galvanizing is the process of cover- ing iron or steel with thin layer of Zn 2. Zinc protects iron sacrificially 3. Zinc protecte iron even when coating of zinc is Punctured 4. Since zinc is toxic galvanized uten- sils are not used for storing food Tinning 1. In tinning, steel is covered with a thin coat of tin to prevent corrosion 2. Due to noble nature tin protects base metal 3. If the coating is puncture intense corrosion of steel occurs. 4. Since tin is non-toxic tinned contain- ers can be used even for storing food (c) Metal cladding [7] Metal cladding is a process of sandwitching a metal between two corrosion resistant metals. Ni, Pb, Cu, Al, Ag, Pt are used as cladding metals over a base metal. For example, aluminium acts as cladding metal for duralumin. The picture of Metal cladding is given in Fig Fig Metal cladding

31 (d) Metal spraying: Molten metal is sprayed on the cleaned surface of a metal for its protection from corrosion. It is helpful to protect towers and bridges (e)Pack cementation: Heating a base metal with another powdered metal is done. The powdered metal forms a coat on the base metal, and hence protects it from corrosion. Sherardising and colourising are two important pack cementation processes in which Zn and Al powders are heated, respectively. This type of coating is also known as diffusion coating. The third one is chromising in which the chromium powder and alumina are heated with Fe (f)Electroplating: It is an electrochemical process in which a base metal is coated by Zn, Ag, Cr, Au, Sn, etc. to protect it from corrosion and also to make it shining and decorative. The base metal is made cathode, dipped in a suitable electrolyte, and the metal to be deposited is made the anode. Anodised coating is done for non-ferrous metals, such as Zn, Mg and Al (g) In organic coatings, paints, varnishes, enamels, lacquers and emulsion paints are included. They not only protect from corrosion but also give a good look to the metal. Organic coats must have chemical inertness, good surface adhesiveness and noneffectiveness towards inorganic chemicals and water (h) Water-repellent paints are organosilicon compounds, which on hydrolysis give a permanent coat that canwork for 4 6 years as corrosion resistant Cathodic protection or electrochemical method: The principle involved in this method is that the metalis forced to behave like cathode, thereby preventing corrosion. For protecting corrosion electrochemically,there are two methods for the protection from corrosion (a) Sacrificial anodic protection. [7] 1. Sacrificial anodic protection: Underground steel pipes are protected from corrosion by this method. A magnesium rod is fixed near the metal under protection and both are connected with a conducting wire. Magnesium is more positive than iron and, hence, in electrochemical cell it acts as anode and the iron acts as cathode. According to the principle of galvanic cell, it is anode that

32 34 undergoes oxidation, and, hence, corrosion occurs at anode saving cathode (iron) from corrosion. Thus, magnesium sacrifies itself for saving the iron. The picture of Sacrificial anodic protection is given in Fig Fig. 1.24Sacrificial anodic protection (b) Impressed current cathodic protection The object to be protected is made cathode and it is connected to the negative terminal of a DC (direct current) source. The positive terminal of the source is connected to the other electrode made of graphite or platinum, lead or nickel. The impressed current opposes the galvanic current (corrosion current) and, hence, protection from corrosion takes place. The picture of Impressed current cathodic protection is given in Fig Fig Impressed current cathodic protection

33 Electrochemical corrosion measurements [44-52] Most metal corrosion occurs via electrochemical reactions at the interface between the metal and an electrolyte solution. A thin film of moisture on a metal surface forms the electrolyte for atmospheric corrosion. Wet concrete is the electrolyte for reinforcing rod corrosion in bridges. Corrosion normally occurs at a rate determined by an equilibrium between opposing electrochemical reactions. The first is the anodic reaction, in which a metal is oxidized, releasing electrons into the metal. The other is the cathodic reaction, in which a solution species (often O 2 or H + ) is reduced, removing electrons from the metal. When these two reactions are in equilibrium, the flow of electrons from each reaction is balanced, and no net electron flow (electrical current) occurs. Corrosion Process Showing Anodic and Cathodic Current Componentsis given in Fig The vertical axis is potential and the horizontal axis is the logarithm of absolute current. The sharp point in the curve is actually the point where the current changes signs as the reaction changes from anodic to cathodic, or vice versa.the sharp point is due to the use of a logarithmic axis. The use of a log axis is necessary because of the wide range of current values that must be displayed during a corrosion experiment. Because of the phenomenon of passivity, it is not uncommon for the current to change by six orders of magnitude during a corrosion experiment. Figure Corrosion Process Showing Anodic and Cathodic Current Components.

34 The potential of the metal The potential of the metal is the means by which the anodic and cathodic reactions are kept in balance. The current from each half reaction depends on the electrochemical potential of the metal. Suppose the anodic reaction releases too many electrons into the metal. Excess electrons shift the potential of the metal more negative, which slows the anodic reaction and speeds up the cathodic reaction. This counteracts the initial perturbation of the system Open Circuit Potential The equilibrium potential assumed by the metal in the absence of electrical connections to the metal is called the Open Circuit Potential, E oc. In most electrochemical corrosion experiments, the first step is the measurement of E oc Corrosion Current The value of either the anodic or cathodic current at E oc is called the Corrosion Current, Icorr. If we could measure I corr, we could use it to calculate the corrosion rate of the metal. Unfortunately, I cor r cannot be measured directly. However, it can be estimated using electrochemical techniques. In any real system, I corr and Corrosion Rate are a function of many system variables including type of metal, solution composition,temperature, solution movement, metal history, and many others. Many metals form an oxide layer on their surface as they corrode. If the oxide layer inhibits further corrosion, the metal is said to passivate. In some cases, local areas of the passive film break down allowing significant metal corrosion to occur in a small area. This phenomena is called pitting corrosion or simply pitting. Because corrosion occurs via electrochemical reactions, electrochemical techniques are ideal for the study of the corrosion processes. In electrochemical studies, a metal sample with a surface area of a few square centimetres 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

35 37 studied. Additional electrodes are immersed in the solution, and all the electrodes are connected to a device called a potentiostat. A potentiostat is used to change the potential of the metal sample in a controlled manner and measure the current the flows as a function of potential. Both controlled potential (potentiostatic) and controlled current (galvanostatic) polarization is useful. When the polarization is done potentiostatically, current is measured, and when it is done galvanostatically, potential is measured. Potentiostatic mode is used to perturb the equilibrium corrosion process. When the potential of a metal sample in solution is forced away from Eoc, it is referred to as polarizing the sample. The response (current) of the metal sample is measured as it is polarized. The response is used to develop a model of the sample's corrosion behaviour. Suppose we use the potentiostat to force the potential to an anodic region (towards positive potentials from Eoc). In Figure 1.26, we are moving towards the top of the graph. This will increase the rate of the anodic reaction (corrosion) and decrease the rate of the cathodic reaction. Since the anodic and cathodic reactions are no longer balanced, a net current will flow from the electronic circuit into the metal sample. The sign of this current is positive by convention. If we take the potential far enough from E oc, the current from the cathodic reaction will be negligible, and the measured current will be a measure of the anodic reaction alone. In Figure 1.26, notice that the curves for the cell current and the anodic current lie on top of each other at very positive potentials. Conversely, at strongly negative potentials, cathodic current dominates the cell current. I corr cannot be measured directly. In many cases, you can estimate it from current versus voltage data. At Eoc, each exponential termequals one. The cell current is therefore zero, Near E oc both exponential terms contribute to the overall current. Finally, as the potential is driven far from Eoc by the potentiostat, oneexponential term predominates and the other term can be ignored. When this occurs, a plot of log currentversus potential becomes a straight line.

36 1.21. Quantitative Corrosion Theory 38 Icorr cannot be measured directly. In many cases, we can estimate it from current versus voltage data. we can measure a log current versus potential curve over a range of about one half volt. The voltage scan is centered on Eoc. we then fit the measured data to a theoretical model of the corrosion process. The model we will use for the corrosion process assumes that the rates of both the anodic and cathodic processes are controlled by the kinetics of the electron transfer reaction at the metal surface. This is generally the case for corrosion reactions. An electrochemical reaction under kinetic control obeys Equation 1-17, the Tafel Equation I = I 0 e(2.3(e-e )/ ) 1.17 where, I = the current resulting from the reaction I 0 = a reaction dependent constant called the Exchange Current E = the electrode potential Eo= the equilibrium potential (constant for a givenreaction) = the reaction's Tafel Constant (constant for agiven reaction). Beta has units of volts/decade. The Tafel equation describes the behaviour of oneisolated reaction. In a corrosion system, we have two opposing reactions anodic and cathodic.the Tafel equations for both the anodic and cathodic reactions in a corrosion system can be combined to generate the Butler-Volmer Equation (Equation 1-18). I = Ia + Ic = Icorr(e(2.3(E-Eoc)/ a) e(-2.3(e-eoc)/ c)) 1.18 where, I = the measured cell current in amps Icorr = the corrosion current in amps E = the electrode potential Eoc = the corrosion potential in volts a = the anodic Beta Tafel Constant in volts/decade c = the cathodic Beta Tafel Constant in volts/decade