UNIT-III: ELECTRO CHEMISTRY AND CORROSION

Transcription

1 UNIT-III: ELECTRO CHEMISTRY AND CORROSION Introduction to Electrochemistry - EMF of the cell or Cell potential-electrochemical series and its importance Reference electrodes (SHE and Calomel electrode). Corrosion (chemical and electrochemical theory of corrosion) Galvanic series. Factors effecting the rate of corrosion Controlling of corrosion (Proper designing, Modifying the environment, Cathodic protections Sacrificial Anodic and Impressed Current Cathodic Protection) 1

2 Electrode potential: The magnitude of the electrode potential of a metal is a measure of its relative tendency to lose or gain electrons, i.e., it is a measure of the relative tendency to undergo oxidation (loss of electrons) or reduction (gain of electrons). It is not possible to measure the absolute value of the single electrode potential directly. Only the difference in potential between two electrodes can be measured experimentally. It is, therefore, necessary to couple the electrode with another electrode whose potential is known. This electrode is termed as reference electrode. The emf of the resulting cell is measured experimentally. The emf of the cell is equal to the sum of potentials on the two electrodes. Emf of the cell = E Anode + E Cathode = Oxidation potential of an + Reduction potential of cathode Knowing the value of reference electrode, the value of other electrode can be determined. Standard electrode potential: The potential difference developed between metal electrode and the solution of its ions of unit molarity (1M) at 25 C (298 K) is called standard electrode potential. According to the IUPAC convention, the reduction potential alone be called as the electrode potential (E O ), i.e., the given value of electrode potential be regarded as reduction potential unless it is specifically mentioned that it is oxidation potential. Standard reduction potential = - (Standard oxidation potential) (Or) Standard oxidation potential = - (Standard reduction potential) emf or cell potential: The difference in potentials of the two half-cells is known as the electromotive force (emf) of the cell or cell potential. The emf of the cell or cell potential can be calculated from the values of electrode potentials of the two half-cells constituting the cell. The following three methods are in use: 1. When oxidation potential of anode and reduction potential of cathode are taken into account: o E Cell = Oxidation potential of anode + Reduction potential of cathode = E o ox (anode) + E o red (cathode) 2. When reduction potentials of both electrodes are taken into account: E o Cell = Reduction potential of cathode - Reduction potential of anode = E o o Cathode E Anode = E o right E left o 3. When oxidation potentials of both electrodes are taken into account: = Oxidation potential of anode - Oxidation potential of cathode E o Cell = E o ox (anode) E o red (cathode) Electrochemical Series: By measuring the potentials of various electrodes versus standard hydrogen electrode (SHE), a series of standard electrode potentials has been established. When elements are arranged in increasing order of their standard electrode potentials, then the series is called electro chemical series. 2

3 Characteristics and applications of electrochemical series: 1) Anode: The negative sign of standard reduction potential indicates that an electrode when joined with SHE acts as anode and oxidation occurs on this electrode. 2) Cathode: The +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode. 3) Reducing agents: The substances which are stronger reducing agents than hydrogen are placed above hydrogen in the series and have negative values of standard reduction potentials. 4) Oxidizing agents: The substances which are stronger oxidizing agents than H + ion are placed below hydrogen in the series. 5) Active metal: The metals on the top (having high negative values of standard reduction potentials) have the tendency to lose electrons readily. These are active metals. 6) Active Non metal: The non-metals on the bottom (having high positive values of standard reduction potentials) have the tendency to accept electrons readily. These are active nonmetals 7) Displacement reactions: A metal higher in the series will displace the metal from its solution of metal is lower in the series, i.e., the metal having low standard reduction potential will displace the metal from its salt's solution which has higher value of standard reduction potential. 8) Feasibility of Redox reaction: Spontaneity of a redox reaction can be predicted from the emf (E) value of the complete cell reaction. Positive value of E of a cell reaction indicates that the reaction is spontaneous. If value of E is negative, the reaction is not feasible. 9) Displacement of hydrogen from water: The metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series tendency to liberate hydrogen gas from dilute acids decreases. 10) Calculating equilibrium constant: Equilibrium constant can be calculated by using the following equation. E 0 = RT/nF ln K eq (Or) E 0 = 2.303RT/ nf log K eq 3

4 Reference Electrodes: (Standard Hydrogen Electrode, SHE or NHE) Hydrogen electrode is the primary standard electrode. It consists of a small platinum strip coated with platinum black as to adsorb hydrogen gas. A platinum wire is welded to the platinum strip and sealed in a glass tube as to make contact with the outer circuit through mercury. The platinum strip and glass tube is surrounded by an outer glass tube which has an inlet for hydrogen gas at the top and a number of holes at the base for the escape of excess of hydrogen gas. The platinum strip is placed in an acid solution which has H + ion concentration 1 M. pure hydrogen gas is circulated at one atmospheric pressure. A part of the gas is adsorbed and the rest escapes through holes. The temperature of the cell is maintained at 25 0 C. By international agreement the standard hydrogen electrode is arbitrarily assigned a potential of exactly ± Volt. The hydrogen electrode thus obtained forms one of two half-cells of a voltaic cell. When this half-cell is connected with any other half-cell, a voltaic cell is constituted. The hydrogen electrode can act as cathode or anode with respect to other electrode. SHE half reaction Electrode potential H 2 2H + + 2e V (Anode) 2H + + 2e - H V (Cathode) Some other reference electrodes: Since a standard hydrogen electrode is difficult to prepare and maintain, it is usually replaced by other reference electrodes, which are known as secondary reference electrodes. These are convenient to handle and are prepared easily. Two important secondary reference electrodes are described here. Calomel electrode: It consists of mercury at the bottom over which a paste of mercury-mercurous chloride is placed. A solution of potassium chloride is then placed over the paste. A platinum wire sealed in a glass tube helps in making the electrical contact. The electrode is connected with the help of the side tube on the left through a salt bridge with the other electrode to make a complete cell. The potential of the calomel electrode depends upon the concentration of the potassium chloride solution. If potassium chloride solution is saturated, the electrode is known as saturated calomel electrode (SCE) and if the potassium chloride solution is 1 N, the electrode is known as normal calomel electrode (NCE) while for 0.1 N potassium chloride solution, the electrode is referred to as decinormal calomel electrode (DNCE). The electrode reactions when the electrode acts as anode and cathode are: 2Hg Hg e - (Oxidation) Hg 2 Cl 2 Hg Cl - Hg Cl - Hg 2 Cl 2 Hg e - 2Hg (Reduction) Hg + 2Cl - Hg 2 Cl 2 + 2e - Hg 2 Cl 2 + 2e - 2Hg + 2Cl The reduction potentials of the calomel electrodes on hydrogen scale at 298K are as follows: Saturated KC V; 1N KC V; 0.1N KC V 4

5 6. J.sureshkumar CORROSION Definition: Any process of deterioration or destruction and consequent loss of a solid metallic material, through in chemical or electro chemical attack by its environment, starting at its surface is called corrosion. Thus, corrosion is reverse of extraction of metals or metallurgy. Corrosion Metal + oxide Metal oxide (Higher Energy) Examples: 1. The most familiar example of corrosion is rusting of iron (Fe 2 O 3. 3H 2 O) 2. Formation of green film of basic carbonate on the surface of copper [CuCO 3 + Cu(OH) 2 ] Causes of Corrosion: Metal exit nature in the form of oxides, sulphides, sulphates and carbonates. These chemically combined states of metals known as mineral or ore has low energy and thermodynamically stable state for the metal. A considerable amount of energy is required for extraction of metal from it ore. The extracted metal has high energy and thermodynamically unstable state. Thus it is the naturally tendency of a metal to back to the thermodynamically stable state. Metal do this by interacting chemically or electrochemically with their environment is known as corrosion. Effects of Corrosion: Effects of corrosion briefly given below. 1. Loss of useful properties of metals and thus loss of efficiency. 2. Decreasing production rate 3. Efficiency of machines decreases by corrosion 4. Increase in maintenance and production cost. 5. Contamination products Types of Mechanism of Corrosion: The mechanism of corrosion classify as two types 1. Dry (or) chemical Corrosion: This type of corrosion occurs mainly through the direct chemical action of environment/atmospheric gases such as-oxygen, halogen, hydrogen sulphide, sulphur dioxide, nitrogen or anhydrous inorganic liquid with metal surfaces immediate proximity. Oxidation Corrosion: It is brought about by the direct action of oxygen on metals usually in the absence of moisture. Alkali metals (Li, Na, K, Rb, etc.) and alkaline-earths (Be, Ca, Sr, etc.) are even rapidly oxidized at low temperatures. At high temperatures, almost all (except Ag, Au, and Pt) are oxidized. The reactions in the oxidation corrosion are: M M e - (Loss of electrons) ½ O 2 + 2e - O 2- (Gain of electrons) M + ½O 2 M 2+ + O 2- (OR) MO (Formation of metal oxide) 5

6 Mechanism: Oxidation occurs first at the surface of the metal and the resulting metal scale forms a barrier that tends to restrict further oxidation. For oxidation to continue either metal must diffuse outwards through the scale to the surface or the oxygen must diffuse inwards through the scale to the underlying metal. Both transfers occur, but the outward diffusion metal is generally much more rapid than the inward diffusion of oxygen, since the metal ion appreciably smaller than the oxygen ion and consequently of much higher mobility. Nature of the oxide formed plays an important part in oxidation corrosion process. A layer is called film, when its thickness is less than about 300 Ǻ and it s called scale, when its thickness exceeds this value. The following types of films are there: 1. Stable: A stable layer is fine grained in structure and can get adhered tightly to the parent metal surface. Hence, such a layer can be of impermeable nature. Such a film behaves as protective coating in nature thereby protecting the surface. The oxide films on Al, Sn, Pb and Cu etc are stable film. 2. Unstable: The oxide layer formed decomposes back into the metal and oxygen. Metal oxide Metal + Oxygen. Ag, Au and Pt metals form unstable film. The film is unstable therefore do not undergo oxidation corrosion in these metals. 3. Volatile: The oxide layer volatilizes as soon as it is formed. Thereby leaving the original metal surface exposed for further attack. This cause rapid and continuous corrosion, leading to excessive corrosion. Molybdenum oxide (MoO 3 ) is volatile. 2Mo + 3O2 4. Porous: 2MoO 3 (Volatile) Having pores or cracks in the film or scale. In such a case, the atmospheric oxygen has entrée to the original surface of metal through the pores or cracks of the film. Thereby the corrosion continues till the entire metal is completely oxide. Alkali, alkaline earth metals and iron form porous films. converted into its PILLING BEDWORTH RULE: It states that an oxide layer is compact, nonporous as well as protective preventing corrosion if the volume of metallic oxide is equal to or greater in volume to the metal surface. The alkali metals like Li, Na, K and alkaline earth metals like Mg and metals like Fe produce oxide film whose volume is less than the volume of the metal and as a result oxygen can diffuse through the pores of the film producing oxide films continuously. This principle also explained according to specific volume ratio. 1. Smaller the specific volume ratio greater the corrosion. 6

7 2. Volume of oxide volume of metal, thus non-porous film and protective film. 3. Volume of oxide < volume of metal, thus porous film and non protective film. WET OR ELECTROCHEMICAL CORROSION: This type of corrosion occurs, where a conducting liquid is in contact with metal or when two dissimilar metals or alloys are either immersed or dipped partially in a solution. The formation of anodic and cathodic areas or parts in contact with each other. Presence a conducting medium. Corrosion takes place at anodic areas only. Formation of corrosion product somewhere between anodic and cathodic areas. Mechanism of wet or electrochemical corrosion: Electrochemical corrosion involves flow of electron-current between the anodic and cathodic. The anodic reaction involves in dissolution of metal as corresponding metallic ions the liberation of free electrons. A cathode reaction consumes electrons with either by two ways: 1. Evolution of Hydrogen Type Corrosion (In acidic environments): Considering metal like Fe, the anodic reaction is dissolution of iron as ferrous ions with the liberation of electrons. At Anode: Fe Fe e (Oxidation) At cathode: 2H + + 2e - H 2 (Reduction) ½ O 2 + 2e - + H 2 O 2 OH - Fe OH Fe (OH) 2 4Fe (OH) 2 + O 2 + 2H 2 O 4 Fe (OH) 3 2Fe (OH) 3 Fe 2 O 3.xH 2 O(Yellow rust) + (3-x) H 2 O (Or) 2Fe (OH) 2 +½ O 2 Fe 3 O 4 (Black rust) + 3H 2 O Thus, this type of corrosion causes displacement of hydrogen ions from the acidic solution by metal ions. Consequently, all metals above hydrogen in the electrochemical series has a tendency to get dissolved in acidic solution with simultaneous 2. Absorption of Oxygen Type Corrosion (In neutral aqueous solution): The surface of iron is, usually coated with at thin film of iron oxide. However, if this iron oxide film develops some cracks, anodic areas are created on the surface; while the well-metal parts act as cathodes. It follows that the anodic areas are small surface parts; while nearly the rest of the surface of the metal forms large cathodes. At Anode: Fe Fe e (Oxidation) At cathode: ½ O 2 + 2e - + H 2 O 2 OH - Fe OH - Fe (OH) 2 \ 4Fe (OH) 2 + O 2 + 2H 2 O Fe (OH) 3 2Fe (OH) 3 (Or) 2Fe (OH) 2 +½ O 2 Fe 2 O 3.xH 2 O (Yellow rust) + (3-x) H 2 O Fe 3 O 4 (Black rust) + 3H 2 O 7 5

8 Differences between Chemical and Electro Chemical Corrosion S.N Chemical Corrosion Electro Chemical Corrosion O 1 It occurs in dry condition It occurs in wet condition (presence of moisture or electrolyte) 2 It involves direct chemical attack of the metal by environment It involves setting up of a large number of galvanic cells 3 It is explained by absorption mechanism It is explained by mechanism of electrochemical reactions 4 It occurs in both homogeneous and It occurs only on heterogeneous surface heterogeneous surfaces 5 Corrosion is uniform Corrosion is not uniform 6 It is a slow process It is a fast process 7 Corrosion product accumulate at the same place where corrosion occurs Corrosion occurs at the anode but products accumulate near cathode GALVANIC SERIES: Electrochemical series give very useful information regarding chemical activity of metals. It did not provide sufficient information regarding corrosion behavior of metals and alloys in environmental conditions. Hence, oxidation potentials of various metals and alloys measured by using standard electrodes immersing the metals and alloys in sea water. Those electrode potentials are arranged in the decreasing order. This series is called galvanic series which give more practical information about the corrosion tendency of various metals and alloys. In this series from top to bottom corrosion tendency decreased i.e. top metals are anodic nature and bottom metals are cathodic nature. Thus it is clear that corrosion occurs at anode part and cathode part to be protected. FACTORS INFLUENCING CORROSION: The rate and extent of corrosion, depends on the following factors: 1. Nature of the metal: (i) Position in galvanic series: When two metals or alloys are in electrical contact, in presence of an electrolyte, the more active metal (or higher up in the series) undergo corrosion. The rate of corrosion depends upon the difference in their positions, and greater is the difference the faster is the corrosion of the anodic metal. (ii) Relative areas of the anodic and cathodic parts: When two dissimilar metals or alloys are in contact, the corrosion of the anodic part is directly proportional to the ratio of areas of the cathodic part and the anodic part. 8

9 (iii) Purity of metal: Impurities in a metal, generally, and form Minute or tiny electrochemical cells and the anodic parts get corroded. For example, zinc metal containing impurity (such as Pb or Fe) undergoes corrosion of zinc, due to the formation of local electrochemical cells. (iv) Physical state of metal: The rate of corrosion is influenced by physical state of the metal.the smaller the grain-size of the metal or alloy, the greater will be its solubility and hence, greater will be its corrosion. (v) Nature of surface film: The ratio of the volumes of the metal oxide to the metal is known as a specific volume ratio. Greater the specific volume ratio, lesser is the oxidation corrosion rate. According to Pilling- bedworth rule the volume of oxide film is greater than metal from which metal oxide formed, then the film is protective. (vi) Solubility of corrosion products: In electrochemical corrosion, if the corrosion product is soluble in the corroding medium, then corrosion proceeds at a faster rate. On the contrary, if the corrosion product is insoluble in the medium or thereby suppressing further corrosion. (vii) Volatility of corrosion products: If the corrosion product is volatile (MoO 3 ), it volatilizes as soon as it is formed, thereby leaving the underlying metal surface ex posed for flirt her attack. This causes rapid and continuous corrosion. (viii) Passivity of metal: Metals like Ti, Al, Cr and Ni are passive they exhibit much higher corrosion resistance, due to the formation of highly protective on the metal. Moreover the film is self healing nature. Thus corrosion resistance of stainless steel is due to passive character of Cr. 2. Nature of the corroding environment: (i) Temperature: With increase of temperature of environment, the reaction as well as diffusion rate increase, thereby corrosion rate is generally enhanced. (ii) Humidity of air: Critical humidity is defined as the relative humidity above which the atmospheric corrosion rate of metal increases sharply. The reason why corrosion of a metal becomes foster in humidity. Atmosphere is that gases (CO 2, O2, etc.) and vapors, present in atmosphere furnish water to the electrolyte, essential for setting up an electrochemical corrosion cell. (iii) Presence of impurities in atmosphere: Atmosphere, in the vicinity of industrial areas, contains corrosive gases like CO 2, H 2 SO 4 and fumes of HCI, H 2 etc. in presence of these gases, the acidity of the liquid, adjacent to the metal surfaces, increases and its electrical conductivity also increases. This consequently, results in an increase of corrosion. (iv) Influence of ph: Generally, acidic media (i.e., ph <7) are more corrosive than alkaline and neutral media. (v) Nature of ions present: Presence of anions like silicate in the medium leads to the formation of insoluble reaction products (e.g., silica gel), which inhibit further corrosion. On the other hand, chloride ions, if present in the medium, destroy the protective and passive surface film. 9

10 Corrosion Control or Protection against Corrosion: Some of the corrosion control methods are described as follows: 1. Proper designing: The design of metal 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 corroding environment. When contact two dissimilar metals, anode should be large are. They should be as close as possible in electrochemical series. Insulating fitting of two dissimilar metals. Anode should not be painted or coated. Welding rather than bolting. Avoid Sharpe corners i.e., Smooth corner rather than sharp corners. 2. Using high purity of metal. 3. Using metal alloys. 4. Modifying the Environment: The corrosive nature of the environment can be reduced either: (i) by the removal of harmful constituents, or (ii) by the addition of specific substances, which neutralize the effect of corrosive constituents of the environment. Deaeration: In oxygen concentration type of corrosion, elimination of oxygen from aqueous environment reduces metal corrosion. The method also reduces the CO 2 content of water, thereby decreasing the corrosion rate. Deactivation: It involves the addition of chemicals, capable of combining rapidly with the oxygen in aqueous solution. For example, sodium sulphite (Na 2 S0 3 ). 2 Na 2S > 2 Na 2S0 4 Hydrazine is useful over the sodium sulphite, because the reaction products are N 2 (g) and water. N 2H >N 2 + 2H 20 Dehumidification: It reduces the moisture content of air to such an extent that the amount of water reduced on metal is too small to cause corrosion. Alumina or silica gels, which adsorbs moisture preferentially on their surfaces. Alkaline neutralization: It is prevention of corrosion by neutralizing the acidic character of corrosive environment due to the presence of acids like H 2 SO 4, HCl, CO 2, S0 2, etc. by alkali like NH 3 NaOH,, lime,etc.) are, generally, injected either in vapour or liquid form to the corroding to its parts. 5. Use of Inhibitors: A corrosion inhibitor is "a substance which when added in small quantities to the aqueous corrosive environment effectively decreases the corrosion of a metal. Inhibitors are two types: Anodic inhibitors: chromates, phosphates, tungstates or other ions of transition elements with high oxygen content are those that suppress the corrosion reaction. They are adsorbed on the metal surface, forming a protective film, thereby reducing the corrosion rate. Cathodic inhibitors: Corrosion may be reduced by slowing down the diffusion 10

11 of hydrated H + ions to cathode. The diffusion of H + ions is considerably decreased by organic inhibitors like amines, heterocyclic nitrogen compounds, substituted ureas and thioureas, heavy metal soaps, Antimony and arsenic oxides. 6. Cathodic protection: The principle involved in this method is to force the metal to be protected to behave like a cathode, thereby corrosion does not occur. There are two types of cathodic protections. (i) Sacrificial anodic protection method: In this protection method, the metallic structure (to be protected) is connected by a wire to a more anodic metal, so that all the corrosion is concentrated at this more active metal. The more active metal itself gets corroded slowly; while the parent structure (cathodic) is protected. The more active metal soemployed is called sacrificial anode. The corroded sacrificial anode block is replaced by a fresh one, when consumed completely. Metals commonly employed as sacrificial anodes are magnesium, zinc, aluminum and their alloys. Important applications of sacrificial anodic method include protection of buried pipelines, underground cables; marine structures, ship-hulls, water-tanks, piers, etc. (ii) Impressed current cathodic protection: In this method, an impressed current is applied in opposite direction to nullify the corrosion current, and convert the corroding metal from anode to cathode. Usually, the impressed current is derived from a direct current source (like battery or rectifier on a.c. line) with an insoluble anode (like graphite, high silica iron, scrap iron, stainless steel or platinum). This type of cathodic protection has been applied to open water-box coolers, water-tanks, buried oil or water pipes, condensers, transmission line towers, marine pier, laid- up ships, etc. 11