Unit-2 Batterychemistry And Corrosion.pdf

UNIT-1 BATERRYCHEMISTRY AND CORROSION BATTERYCHEMISTRY  Battery is an electrochemical cell or often several electrochemical cells connected in series that can be used as a source...

UNIT-1 BATERRYCHEMISTRY AND CORROSION BATTERYCHEMISTRY  Battery is an electrochemical cell or often several electrochemical cells connected in series that can be used as a source of direct electric current at a constant voltage.  A device which converts chemical energy to electrical energy is called battery.  Usually in the term applied to a group or two or more electric cells, connected together electrically in series.  Batteries are commercial electrochemical cells.  Batteries are 3 types 1. Primary cells 2. Secondary cells 3. Fuel cells Primary cells:  These are disposable batteries, which are designed to use once and discard.  Cannot be recharged  The cell reaction is not reversible; hence when all the reactants are converted into products, no more electricity is produced and the cell becomes dead and cannot be used after that. (the reactants are not return back to reactants by application of current)  These batteries are used as a source of DC power.  These cells are convenient to use and inexpensive.  Used in ordinary gadgets like torch lights, watches and toys.   Ex: Leclanche cell, Lithium cells and alkaline cells. Secondary cells / Storage cells:  These are rechargeable batteries, which are designed to be recharged and used multiple times.  The cells in which the cell reaction is reversed by passing direct current in opposite direction.  The secondary batteries can be used through a large number of cycles of discharging and charging.  They are used as source of dc power used to supply large, short term repetitive power requirements such as automotive and airplane batteries.  These batteries have very large capacitance and ling periods of low current rate discharge.  The cell reaction is reversible; the reactants are return back to reactants by application of current  These batteries are used as a source of DC power.  EX: Lead acid storage cell (Lead accumulator) Lead accumulator or Lead storage battery (Pb-acid battery):  A storage cell can operate both as a voltaic cell and as an electrolytic cell.  When operating as a voltaic cell, it supplies electrical energy and as a result eventually becomes rundown, when being recharged; the cell operates as an electrolytic cell.  This storage cell has the great advantage of working both as an electrolytic cell and as a voltaic cell.  It has lead anode and PbO2 cathode and dilute H2SO4 acts as electrolyte.  These all are placed in a plastic container as shown in diagram.  A number of lead plates (-ve plates) are connected in parallel and a number of lead dioxide plates (+ve plates) are also connected in parallel.  The lead plates fit in between the lead dioxide plates. The plates are separated from adjacent one by insulators like wooden strips, rubber or glass fibre.  The entire combination is immersed in 20% dilute H2SO4 corresponsing specific gravity of 1.2  During Discharge : o In the discharged state:  Both the anode is lead (Pb) and cathode is leadoxide / lead peroxide (PbO2) become lead (II) sulfate (PbSO4) and the dilute H2SO4 is changed into water.  Electrons move from cathode to anode through bulb.  The cell reactions are At anode: Pb (s) + H2SO4 (aq) → PbSO4 (s) + 2H+ (aq) + 2e At cathode: PbO2 (s) + H2SO4 (aq) + 2H+ (aq) + 2e → PbSO4 (s) + 2H2O (l) Net reaction: Pb (s) + PbO2 (s) + 2H2SO4 (aq) → 2PbSO4 (s) + 2H2O (l) + energy  During Charging : o In the charged state:  Cell has anode of leadoxide (PbO2) and cathode of lead (Pb) and dilute H2SO4  Electrons move from anode to cathode.  The cell reactions are At anode: PbSO4(s) + 2H+ (aq) + 2e → Pb(s) + H2SO4 (aq) At cathode: PbSO4(s) + 2H2O (l) → PbO2(s) + H2SO4 (aq) + 2H+ (aq) + 2e Net reaction: 2 PbSO4 (S) + 2H2O (l) → Pb (s) + PbO2 (s) + 2 H2SO4 (aq) OR  Applications: o Lead storage batteries are used in cars, motor cycles, Lorries to start the engine. o They are also used in invertors for electric supply in telephone exchanges, railway trains, hospitals, automobiles, power-stations and houses. Advantages: 1. It has relatively constant potential i.e. 12V 2. It is portable and inexpensive. 3. When electricity is being drawn from the cell, to start the car it acts as a voltaic cell, when the car is running, the cell is being recharged as an electrolytic cell. Lithium-ion battery:  Fuel cell  A fuel cell is an electrochemical cell which converts chemical energy contained in an easily available fuel oxidant system into electrical energy. The basic principles of fuel cells are identical to those of the electrochemical cells, the only difference is that in fuel cell the chemical energy is provided by a fuel and an oxidant stored outside the cell.  A fuel cell changes the chemical energy of the fuel into electricity. Since the electrons exchange takes place when a fuel gets oxidized.  They are of 2 types.Those are:  (1) H2-O2 Fuel cell and   (2) methyl alcohol-oxygen fuel cell.  The essential process in fuel cell is Fuel + oxygen → oxidation products + electricity  Representation of fuel cell: Fuel │electrode │ electrolyte │ electrode │ oxidizer  In this cell the fuel and oxidizer are continuously supplied separately to the electrodes of the cell.  Fuel cells are capable of supplying current as long as the fuel and oxidizer are supplied.  Principle: o The catalyst present in the anode pull off electrons from the H-rich fuel at the anode to produce H+ ions then this H+ ions travel through the electrolyte from anode to cathode. o Electrons formed at the anode travel through the circuit, give direct current and come to the cathode. o There the electrons, H+ ions and O2 form water and this water comes out from the fuel cell as shown in the diagram. H2-O2 Fuel cell:  It has as an electrolyte (2.5% KOH) and two inert electrodes having holes.  H2 and O2 gases are passed through the anode and cathode rooms respectively.  The electrodes made up of graphite are saturated with finely powdered platinum.  A number of such fuel cells are joined together in series to make a battery.  The following reactions take place in anode and cathode rooms are: At anode: 2H2 g + 4 OH − (aq) → 4H2 O l + 4 e At cathode: O2 g + 2H2 O l + 4e → 4 OH − (aq) Net reaction: 2H2 g + O2 g → 2H2 O (l) 0 0 o The standard EMF of the cell, E 0 = Eoxd + Ered = 0.83 + 0.40 = 1.23 V o Applications: o H2-O2 fuel cells are used as energy sources in space vehicles, submarines & military vehicles. o 250 kg of H2 + 02 is enough to produce electricity for 15 days  o The water formed from fuel cell is very useful to astronauts. Advantages:  The energy conversion is very high (75-82%)  Noise and thermal pollution are low.  The maintenance cost is low for these fuels.  The regenerative hydrogen – oxygen fuel cell is an energy storage system for space applications. Limitations:  The life time of fuel cells are not accurately known.  Their initial cost is high.  The distribution of hydrogen is not proper. Solar Energy:  Solar power is the conversion of energy from sunlight into electricity, either directly using photovoltaics (PV), indirectly using concentrated solar power, or a combination.  Concentrated solar power systems use lenses or mirrors and solar tracking systems to focus a large area of sunlight into a small beam. Photovoltaic cells convert light into an electric current using the photovoltaic effect.  Photovoltaics were initially solely used as a source of electricity for small and medium- sized applications, from the calculator powered by a single solar cell to remote homes powered by an off-grid rooftop PV system.  Commercial concentrated solar power plants were first developed in the 1980s.  As the cost of solar electricity has fallen, the number of grid-connected solar PV systems has grown into the millions and utility-scale photovoltaic power stations with hundreds of megawatts are being built.  Solar PV is rapidly becoming an inexpensive, low-carbon technology to harness renewable energy from the Sun.  The current largest photovoltaic power station in the world is the Pavagada Solar Park, Karnataka, India with a generation capacity of 2050 MW. Solar Cell – Photovoltaic Cell:  A solar cell, or photovoltaic cell (PV), is a device that converts light into electric current using the photovoltaic effect.  The first solar cell was constructed by Charles Fritts in the 1880s.  The German industrialist Ernst Werner von Siemens was among those who recognized the importance of this discovery  The solar cell is the elementary building block of the photovoltaic technology. Solar cells are made of semiconductor materials, such as silicon.  One of the properties of semiconductors that makes them most useful is that their conductivity may easily be modified by introducing impurities into their crystal lattice.  There are several types of solar cells. However, more than 90 % of the solar cells currently made worldwide consist of wafer-based silicon cells.  They are cut either from a single crystal rod or from a block composed of many crystals and are correspondingly called mono-crystalline or multi- crystalline silicon solar cells.  Wafer-based silicon solar cells are approximately 200 μm thick. Another important family of solar cells is based on thin-films, which are approximately 1-2 μm thick and therefore require significantly less active, semiconducting material.  Thin-film solar cells can be manufactured at lower cost in large production quantities; hence their market share will likely increase in the future. However, they indicate lower efficiencies than wafer-based silicon solar cells, which means that more exposure surface and material for the installation is required for a similar performance. Principle: The basic principle involved in the solar cells is based on thephotovoltaic (PV) effect. When sun rays fall on the two layers ofsemiconductor devices, potential difference between the two layers isproduced. This potential difference causes flow of electrons and thusproduces electricity. Example: Silicon solar cell. Construction: Solar cell consists of a p-type (such as Si doped with boron) and a n-type (such as Si doped with phosphorous). They are in close contact witheach other. Working: When the solar rays fall on the top layer of p-type semiconductor, the electrons from the valence band get promoted to the conduction band and cross the p-n junction into n-type semiconductor. Thereby potential difference between two layers is created, which causes flow of electrons (i.e. electric current). The potential difference and hence current increases as more solar rays falls on the surface of the top layer. Thus, when this p- and n- layers are connected to an external circuit ,electrons flow from n-layer to p-layer and hence current is generated. Photovoltaic Cell Set-up:  A number of solar cells electrically connected to each other and mounted in a single support structure or frame is called a ‘photovoltaic module’.  Modules are designed to supply electricity at a certain voltage, such as a common 12- volt system. The current produced is directly dependent on the intensity of light reaching the module.  Several modules can be wired together to form an array. Photovoltaic modules and arrays produce direct-current electricity.  They can be connected in both series and parallel electrical arrangements to produce any required voltage and current combination.  There are two main types of photovoltaic system.  Grid- connected systems (on-grid systems) are connected to the grid and inject the electricity into the grid.  For this reason, the direct current produced by the solar modules is converted into a grid-compatible alternating current. However, solar power plants can also be operated without the grid and are then called autonomous systems (off-grid systems). More than 90 % of photovoltaic systems worldwide are currently implemented as grid- connected systems.  The power-conditioning unit also monitors the functioning of the system and the grid and switches off the system in case of faults. On-grid Vs Off-grid:  On-grid means your solar system is tied to your local utility’s GRID. This is what most residential homes will use because you are covered if your solar system under or over-produces in regard to your varying energy needs.  This entire means for you is that your utility system acts as your battery space.  If you are producing more energy with your solar panels or system than you are using, the excess energy is sent to your grid’s power company, allowing you to build credit that you can cash out with at the end of the year, in a process called net metering.  Being grid-tied is beneficial because you do not have to buy an expensive battery back-up system to store any excess energy.  Off-grid means, you are not connected in any way to your grid’s power system or utility company.  This is appealing because you are 100% self-sustaining your energy use.  However, there are disadvantages because off-grid systems require you to purchase back-up battery, which can be expensive, bulky, and not very environmentally friendly, which defeats the purpose of going solar (save money and live greener).  If you do not have batteries or a means to store your energy, you will have less or no electricity when it is cloudy, and you will not have electricity at night.  Applications of solar cells: 1. Solar cells are used in street lights. 2. Water pumps are operated by using solar batteries. 3. They are used in calculators, watches, radios and TVs. 4. They are used for eco-friendly driving vehicles. 5. Silicon Solar cells are used as power source in space crafts and satellites. 6. Solar cells can even be used in remote places and in forests to get electrical energy without affecting the atmosphere. CORROSION Introduction: The surface of almost all the metals begin to decay more or less rapidly when exposed to atmospheric gases, water or other reactive liquid medium.  The process of decay of metal by environmental attack is known as corrosion.  Metals undergo corrosion and convert to their oxides, hydroxides, carbonates, sulphides, etc.  Eg. Iron undergoes corrosion to form reddish brown colour rust [Fe2O3. 3H2O].  Copper undergoes corrosion to form a green film of basic carbonate [CuCO3 + Cu(OH)2] Causes of corrosion: 1. The metals exist in nature in the form of their minerals or ores in the stable combined forms as oxides, chlorides, silicates, carbonates and sulphides. 2. During the extraction of metals, these ores are reduced to metallic state by supplying considerable amount of energy. 3. Hence the isolated pure metals are in excited states than their corresponding ores. So metals have natural tendency to go back to their combined state (minerals/ores). When metal is exposed to atmospheric gases, moisture, liquids etc, the metal surface reacts and forms more thermodynamically stable compounds. Effects of corrosion: 1. Wastage of metal in the form of its compounds. 2. The valuable metallic properties like conductivity, malleability, ductility etc. are lost due to corrosion. 3. Life span and efficiency of metallic parts of machinery and fabrications is reduced Theories of corrosion: Dry corrosion or Chemical corrosion: This type of corrosion occurs mainly through the direct chemical action of atmospheric gases like O2, halogens, H2S, SO2, N2 or anhydrous inorganic liquid with the metal surface. Wet Corrosion or Electrochemical Corrosion  This type of corrosion occurs when a conducting liquid is in contact with the metal. This is due to the existence of separate anodic and cathodic parts, between which current flows through the conducting solution.  At anodic area, oxidation reaction occurs there by destroying the anodic metal either by dissolution or formation of compounds. Hence corrosion always occurs at anodic parts. Mechanism: Electrochemical corrosion involves flow of electrons between anode and cathode. The anodic reaction involves dissolution of metal liberating free electrons. M Mn+ + ne- The cathodic reaction consumes electrons with either evolution of hydrogen or absorption of oxygen which depends on the nature of corrosive environment. Wet corrosion takes place in two ways. 1. Evolution of Hydrogen 2. Absorption of Oxygen  Evolution of Hydrogen: This type of corrosion occurs in acidic medium. Eg: Considering the metal Fe, anodic reaction is dissolution of iron as ferrous ions with liberation of electrons Anode: Fe Fe2+ + 2e- (Oxidation) The electrons released flow through the metal from anode to cathode, whereas H+ ions of acidic solution are eliminated as hydrogen gas. Cathode: 2H+ + 2e- H2 (Reduction) The overall reaction is: Fe + 2H+ Fe2+ + H2 This type of corrosion causes displacement of hydrogen ions from the solution by metal ions. All metals above hydrogen in electrochemical series have a tendency to get dissolved in acidic solution with simultaneous evolution of H2 gas. In this case anodic area is large and cathodic area is small. Absorption of Oxygen: For example, rusting of iron in neutral or basic aqueous solution of electrolytes in presence of atmospheric oxygen. Usually the surface of iron is coated with a thin film of iron oxide. If the film develops cracks, anodic areas are created on the surface and the rest of the metal surface acts as cathodes. It shows that anodes are small areas, while the rest metallic part forms large. Cathode: The released electrons flow from anode to cathode through iron metal. At anode: Fe Fe2+ + 2e- (Oxidation) - At cathode: ½ O2 + H2O + 2e- 2OH (Reduction) - Overall reaction: Fe2+ + 2OH Fe(OH)2  If oxygen is in excess, ferrous hydroxide is easily oxidized to ferric hydroxide. 4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3 4Fe(OH ) O 2H O oxida  tion 4Fe(OH ) oxida  tion Fe O.3H O 2 2 2 3 2 3 2 The product called yellow rust corresponds to Fe2O3.3H2O. Differences between chemical and electrochemical corrosion Types of Corrosion: Galvanic Corrosion:  When two dissimilar metals are electrically connected and exposed to an electrolyte, the metal higher in electrochemical series (low reduction potential) undergoes corrosion and the metal lower in electrochemical series (high reduction potential) is protected. This type of corrosion is called galvanic corrosion.  Eg: When Zn an Cu are connected and exposed to corroding environment, Zinc higher in electrochemical series forms anode, undergoes oxidation and gets corroded. Cu lower in electrochemical series acts as cathode, undergoes reduction and protected as the electrons released by Zn flow towards Cu. Galvanic Series: Although electrochemical series give very useful information regarding chemical reactivity of metals, it did not provide sufficient information in predicting the corrosion behavior under a particular set of environmental conditions. Hence electrode potentials of various metals and their alloys in common use are measured by immersing them partially in sea water, and the values have been arranged in decreasing order of activity and is called as galvanic series. In galvanic series, oxidation potentials are arranged in the decreasing order of activity of a series of metals. Thus galvanic series give real and useful information for studying the corrosion tendency of metals and alloys. Pitting corrosion:  Pitting corrosion is a localized accelerated attack, resulting in the formation of cavities.  It results in the formation of pin holes, pits and cavities on the metal surface due to  cracking of the protective film on a metal at specific points  A cavity, pinholes and cracks on the protective film developed on the metal surface creates the formation of small anodic areas in the less oxygenated parts and large cathodic areas in well oxygenated parts.  The corrosion product is formed between cathodic and anodic areas.  Cracking of protective film may be due to surface roughness or non uniform stresses. Scratches, local strain of the metal, alternating stresses, sliding under load and chemical attack Water line corrosion  This type of corrosion is due to differential aeration corrosion.  When water is stored in a steel tank, it is generally found that the maximum amount of corrosion takes place along a line just below the water line.  This is because the area below water line (poorly oxygenated) acts as anode and gets corroded.  The area above the water line (highly oxygenated) act as cathode and is completely unaffected by corrosion Factors effecting corrosion: The rate and extent of corrosion depends upon various factors due to nature of metal and nature of corroding environment. Nature of metal: 1. Purity of the metal: If impurities are present in a metal the corrosion rate is increased due to formation of tiny electrochemical cells at the exposed parts and the anodic parts get corroded. A pure metal is more corrosion resistant than impure metal. By increasing its purity the corrosion resistance of a metal can be improved Ex: Zn metal containing impurity undergoes corrosion of zinc, due to the formation of electrochemical cells. The rate and extent of corrosion increases with the increasing exposure and extent of the impurities. 2. Electrode potentials: Metals with higher reduction potentials do not corrode easily. They are noble metals like gold, platinum and silver. Whereas the metals with lower reduction potentials readily undergo corrosion (eg. Zn, Mg, Al etc.). 3. Position of metal in galvanic series: Metals which possess low reduction potentials and occupy higher end of galvanic series undergo corrosion easily. 4. Metals which possess high reduction potentials and occupy lower end of galvanic series do not undergo corrosion and they get protected.  When two metals are in electrical contact in presence of an electrolyte, then the metal which is more active undergoes corrosion.  The rate of corrosion depends on the difference in their position in galvanic series. Greater the difference more will be the extent of corrosion at anode. Eg: The potential difference between Fe and Cu is 0.78V which is more than that between Fe and Sn (0.30V). Therefore, Fe corrodes faster when in contact with Cu than that with Sn. On this account, the use of dissimilar metals should be avoided wherever possible (Eg. Bolt & nuts, screw & washer). 5. Relative areas of anodic and cathodic parts: If the metal has small anodic and large cathodic area, the rate of corrosion is very high. This is because the more electrons are liberated at smaller anodic area, which are consumed at cathode. If the cathodic area is larger, the liberated electrons are rapidly consumed at cathode. This further enhances the anodic reaction leading to increase in the rate of corrosion. When two dissimilar metals or alloys are in contact, Area of cathodic part corrosion at anodic area ∝ Areaof anodic part , 6. Hydrogen over voltage: The difference between the potential of the electrode at which the electrolysis actually proceeds continuously (actual decomposition potential) and the theoretical decomposition potential for the same solution is called overvoltage. When a metal which is at a high position in galvanic series (eg. Zn) is placed in 1N H2SO4, it undergoes corrosion with deposition of a film on its surface and evolution of hydrogen gas. So the initial rate of corrosion is high, but decreases after a while due to salt film and H2 film surrounding the metal which causes high over voltage and reduces the corrosion rate. However, if few drops of CuSO4 are added, the corrosion rate of Zn is accelerated because some copper gets deposited on the Zn metal, forming minute cathodes, where the hydrogen overvoltage is reduced. Hence reduction in overvoltage of the corroding metal/alloy accelerates the corrosion. So higher the over voltage, lesser is the corrosion 1 Rate of corrosion α over voltage  Physical state of metal: Metals with small grain size have more tendencies to undergo corrosion. Metal with more stress/strain also undergoes corrosion easily. Nature of surface film: If the corrosion product formed is more stable, insoluble and nonporous, it acts as protective layer and prevents further corrosion (Eg. Ti, Al and Cr). If the corrosion product is porous, volatile and soluble, it further enhances the corrosion (Fe, Zn and Mg). Volatility of corrosion product: If the corrosion produced volatilizes as soon as it is formed the metal surface is exposed for further attack. This creates rapid and excessive corrosion. For example the corrosion product of molybdenum as molybdenum oxide is volatile. Solubility of corrosion product: If the oxide film formed as corrosion product is soluble in corroding medium the corrosion proceeds at a faster rate. The corrosion product acts as a physical barrier between the metal and environment. For example, PbSO4 film formed by Pb in sulphuric acid medium. Nature of Environment: 1. Temperature: The rate of corrosion increases with increase in temperature due to increase in diffusion rate. 2. Humidity in air: The rate of corrosion increases with the presence of moisture in atmosphere because the moisture or humidity present in atmosphere furnishes water to the electrolyte which is essential for setting up of an electrochemical cell. The oxide film formed has the tendency to absorb moisture which creates another electrochemical cell. 3. Presence of impurities: Atmosphere is contaminated with gases like CO2, SO2, H2S; fumes of H2SO4, HCl etc. and other suspended particles in the vicinity of industrial areas. They are responsible for electrical conductivity, thereby increasing corrosion. 4. Effect of PH: pH value of the medium has the greater effect on corrosion. Generally acidic medium (i.e. pH < 7) is more corrosive than basic medium. Acidic pH increases the rate of corrosion. However some metals like Al, Zn, Pb etc dissolve in alkaline solutions as complex ions. Consequently, corrosion of metals, readily attacked by acid can be reduced by increasing the PH of the attacking environment. Acidic medium: PH < 7 - Corrosion is more Basic medium: PH > 7 - Corrosion is less Eg: Zn which is readily corroded in acidic solutions suffers very less corrosion in alkaline medium, i.e PH =11. Al has less corrosion at pH=5.5 which corrodes rapidly at PH = 8.5. 5. Amount of oxygen in atmosphere: As the percentage of oxygen in atmosphere increases, the rate of corrosion also increases due to the formation of oxygen concentration cell. The decay of metal occurs at the anodic part and the cathodic part of the metal is protected.  Corrosion control methods: I. Cathodic Protection: The method of protecting the base metal by making it to behave like a cathode is called as cathodic protection. There are two types of cathodic protection (a) Sacrificial anodic protection (b) Impressed current protection a) Sacrificial anodic protection  In this protection method, the metallic structure to be protected (base metal) is connected by a wire to a more anodic metal so that all the corrosion is concentrated at this more anodic metal.  The more anodic metal itself gets corroded slowly, while the parent structure (cathodic) is protected. The more active metal so employed is called sacrificial anode. The corroded sacrificial anode is replaced by a fresh one, when consumed completely.  Metals commonly employed as sacrificial anode are Mg, Zn, Al and their alloys which possess low reduction potential and occupies higher end in electrochemical series. Eg: A ship-hull which is made up of steel is connected to sacrificial anode (Zn-blocks) which undergoes corrosion leaving the base metal protected. Figure1. Sacrificial anode method: Ship hull and underground water pipeline Applications of Sacrificial anodic protection: By referring to the electrochemical series, the metal with low reduction potential is connected to the base metal which acts as anode. 1. To protect underground pipelines- Buried pipe line protected by connecting to Mg block  2. Protection of ship hulls and other marine devices 3. Protection of water tank- by suspending Zn or Mg rods, body of the tank made cathode and protected 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.  The impressed current is slightly higher than the corrosion current. Thus the anodic corroding metal becomes cathodic and protected from corrosion.  The impressed current is taken from a battery or rectified on A.C. line.  The metal to be protected is made cathode by connecting to an external battery (-ve terminal)  The anode is usually insoluble anode like graphite, high silica iron, scrap iron, stainless steel, or platinum. Usually a sufficient D.C current is passed on to the insoluble anode kept in a black fill composed of coke or gypsum, so as to increase the electrical contact with the surrounding soil.  In impressed current cathodic protection, electrons are supplied from an external cell, so that the object itself becomes cathodic and not oxidized. Applications:  The impressed current cathodic protection is used for the protection of water tanks, water & oil pipe lines, transmission line towers etc. b) Metallic coatings: The surface of the base metal is coated with another metal (coating metal) is called metallic coatings. Metallic coatings are broadly classified into anodic and cathodic coatings. 1. Anodic coating:  The metal used for the surface coating is more anodic than the base metal which is to be protected.  For example, coating of Al, Cd and Zn on steel surface are anodic because their electrode potentials are lower than that of the base metal iron. Therefore, anodic coatings protect the underlying base metal sacrificially.  The formation of pores and cracks over the metallic coating exposes the base metal and a galvanic cell is formed between the base metal and coating metal. The coating metal dissolves anodically and the base metal is protected. 2. Cathodic coating:  Cathodic coatings are obtained by coating a more noble metal (i.e. metals having higher electrode potential like Sn, Au, Ag, Pt etc.) than the base metal. They protect the base metal as they have higher corrosion resistance than the base metal due to cathodic nature.  Cathodic coating protects the base metal only when the coating is uniform and free from pores.  The formation of pores over the cathodic coating exposes the base metal (anode) to environment and a galvanic cell is set up. This causes more damage to the base metal. Methods of application of metallic coatings: 1. Hot dipping:  Hot dipping process is applicable to the metals having higher melting point than the coating metal.  It is carried out by immersing a well cleaned base metal in a bath containing molten coating metal and a flux layer.  The flux cleans the surface of the base metal and prevents the oxidation of the molten coating metal. Eg: Coating of Zn, Pb, Al on iron and steel surfaces. The most widely used hot dipping processes are (a) Galvanizing (b) Tinning Galvanizing:  Galvanizing is a process in which the iron article is protected from corrosion by coating it with a thin layer of zinc.  It is the anodic protection offered by the zinc. In this process, at first iron or steel is cleaned by pickling with dilute sulphuric acid solution at a temperature range of 60-900C for 15 to 20 minutes. Therefore, it removes scale, rust and other impurities present and then washed well in a water bath and dried.  Then after dipped in the bath containing molten zinc which is at 425-450oC. To prevent it from oxide formation, the surface of bath is covered with an ammonium chloride flux. Then the iron sheet is taken out which is coated with a thin layer of zinc.  To remove excess zinc, it is passed through a pair of hot rollers and then it is annealed at a temperature of 450oC followed by cooling.  Applications: Galvanizing is widely used for protecting iron exposed to the atmosphere (roofs, wire fences, pipes etc.) Galvanized metallic sheets are not used for keeping eatables because of the solubility of zinc. Tinning  The process of coating tin over the iron or steel articles to protect them from undergoing corrosion is known as tinning.  Tin is a noble metal and therefore it possess more resistance to chemical attack. It is the cathodic protection offered by the tin. In this process, iron sheet is treated in dilute sulphuric acid (pickling) to remove any oxide film, if present.  A cleaned iron sheet is passed through a bath ZnCl2 molten flux followed by molten tin and finally through a suitable vegetable oil. The ZnCl2 flux helps the molten metal to adhere to the base metallic surface.  Palm oil protects the tin coated surface against oxidation. Applications: 1. Tinning of mild steel plates is done mostly for the requirements of the food stuff industry. 2. Tin metal possess good resistance against atmospheric corrosion. Tin is non-toxic and widely used for coating steel, copper and brass sheets 3. Tinned copper sheets are used for making cooking utensils and refrigeration equipment.