There Are Two Types of Voltaic Cell

CHEMISTRY ASSIGNMENTElva Kumalasari XI-IC The electrolyte includes a mixture of magnesium (IV) oxide and carbon powder. while the negative terminal is the zinc casing around the cell. A galvanic cell is also called a voltaic cell. The picture shows a copper zinc galvanic cell (battery). surrounded by ammonium chloride powder. Two half cells can be put together to form an electrolytic cell. which can be recharge Application of voltaic cell in battery (Dry Cell) Dry Cell (one of secondary cell) Chemistry is the driving force behind the magics of batteries. A galvanic cell consists of at least two half cells. At the positive terminal (carbon) NH4+ ions are discharged. which cannot be recharge Secondary Cell. which results in this reaction. which is used for electrolysis. A battery is a package of one or more galvanic cells used for the production and storage of electric energy by chemical means.dissolve in the electrolyte. The spontaneous reactions in it provide the electric energy or current. ammonia and hydrogen. Chemical reactions in the two half cells provide the energy for the galvanic cell operations. They receive electrons to form two gases. 2NH+4(aq) 2NH3(g) + H2(g) The hydrogen. The positive terminal of a dry cell is a carbon rod. Each half cell consists of an electrode and an electrolyte solution. In this case. electric energy is used to force nonsponaneous chemical reactions. We will make this introduction using a typical setup as depicted here.There are two types of Voltaic Cell: Primary Cell. The chemical reaction which takes place are: At the negative terminal. Zn: Zn (s) Zn2+(aq) + 2eZn2+ ions. a reduction cell and an oxidation cell. - - . reacts with manganese (IV) oxide as follows: 2MnO2(s) + H2(g) Mn2O3(g) + H2O(l) Overall reaction Zn (s) + 2MnO2(s) + 2NH+4(aq) Zn2+(aq) + Mn2O3(s) + 2NH3(g) + H2O(l) Carbon powder is used to increase the surface area of the carbon electrode and manganese (IV) oxide reduces the formation of gas bubbles. which form when Zn donates electrons. Usually the solution contains ions derived from the electrode by oxidation or reduction reaction. . Therefore. plastic or glass fibre. The lead-sulphuric acid cells is a common example. except for some energy loss. Electrodes: Alternate parallel plates of lead dioxide (+ve electrode) (oxidised from PbO) Spongy lead (reduced from PbO) (-ve electrode) These are kept separate by porous separators made of wood. The hydrogen ions go to the +ve electrode and SO4 to the -ve electrode. in 1859. at the -ve electrode 2- At the +ve electrode. Electrolytes: Dilute sulphuric acid Container: Glass or bakelite Here stored chemical energy is converted to electrical energy or current is drawn from the cell. Thus the cell recovers its original state. Such cells are called secondary cells or accumulators. Both plates (but only half of the active materials) are converted into PbSO4 (whitish). Gaston plate. After giving their charges they react with the electrodes and reduce the active material to lead sulphate. It was inverted by a French physicist. Water is formed thus lowering the specific gravity of H2 SO4 (electrolyte). The chemical processes that occur at the electrodes during discharge are reversed by this.Accu (one of primary cell) Lead Accumulator This can be recharged by passing a current through it in the reverse direction. The emf of the cell falls and sulphuric acid is consumed. Water is consumed and sulphuric acid is formed thus raising the specific gravity of the electrolyte. At -ve electrode. the +ve electrode is coated with dark brown lead peroxide and the -ve electrode with grey spongy lead. Both plates (but only half of the active materials) are converted into PbSO4 (whitish). and the cathode to the negative terminal.Discharging Process Here stored chemical energy is converted to electrical energy or current is drawn from the cell.c. The emf of the cell falls and sulphuric acid is consumed. source. Water is formed thus lowering the specific gravity of H2 SO4 (electrolyte). Therefore. After giving their charges they react with the electrodes and reduce the active material to lead sulphate. That is. At the +ve electrode. the anode is connected to the positive terminal of the d. Recharging Process Current is passed through the two terminals in the reverse direction to that in which the cell provided current. . at the -ve electrode 2- At the +ve electrode. and the electrical energy supplied is converted into chemical energy which is stored in the cell. The hydrogen ions go to the +ve electrode and SO4 to the -ve electrode. In the charging process. The emf of the cell rises. The hydrogen ions move to the -ve electrode and sulphate ions to the +ve electrode. The major method for producing it is the electrolysis of brine or "salt water.Application of electrolysis Production of Chemical Example: production of sodium chloride. constitute sodium hydroxide. is one of the most important of all industrial chemicals. migrate to the positive anode and lose their electrons to become chlorine gas. and the lye is used in making soap and paper. Chlorine and hydrogen gases are produced as valuable byproducts. together with the sodium ions that are already in the solution. also known as lye and caustic soda. hydrogen and sodium hydroxide. which can be recovered by evaporation. the chlorine is used in the purification of water. Na +. . Cl -. the hydrogen is used in the hydrogenation of oils. NaOH. (The chlorine atoms then pair up to form Cl 2molecules. Among other uses. But they do not pick up electrons to become sodium metal atoms as they do in molten salt. sodiu m ions. The hydroxide ions. It is produced at the rate of 25 billion pounds a year in the United States alone. the negative chloride ions. chlorine and hydrogen Sodium hydroxide. By electricity. it converts cheap salt into valuable chlorine.) Meanwhile. This so-called chloralkali process is the basis of an industry that has existed for well over a hundred years. When an electric current is passed through salt water. are drawn to the negative cathode." a solution of common salt. sodium chloride in water. because in a water solution the water molecules themselves pick up electrons more easily than sodium ions do. As the oxygen comes in contact with the ionic complex it causes the reduction of the copper (I) ions in the ore to copper metal. are not oxidized at the anode. and gold. The impure copper is further refined via electrolysis. The more electropositive zinc and iron are oxidized into their respective ions and enter into solution. Cu2S + O2 -----> 2 Cu (s) + SO2 (g) The copper metal that is formed by this reduction process still contains a small amount of impurities such as zinc. At the same time sulfur is oxidized to sulfur dioxide as shown in the following equation. In this process. The Noble metals. silver and gold. but settle out as metal atoms in a "sludge" as the impure copper anode dissolves. as shown below. A common copper containing ore is chalcocite (Cu2S). iron.   . roasting) or by blowing oxygen through the melted ore. The copper obtained by this refinement is about 99. silver.5% pure.Purification of metal/ efining metal Example of purification of metal (purifying copper) An example of the process described above is the refinement of copper. the impurities are removed in one of two ways. This ore is first treated at high temperatures (by a process called. and the amount of silver and gold recovered in the process is often sufficient to pay for the cost the electricity required for the electrolytic process. a rectifier. ore commonly. any plating baths include cyanides of other metals e. copper is oxidized at the anode to u2+ y losing two electrons. dditionally. zero valence state.El t oplatin e c hode he The e ec oplating cell are oth connected to an e ternal s pply of direct c rrent . some alloys can e electrodeposited. not an alloy.in the solution to form copper sulfate. The result is the effective transfer of copper from the anode source to a plate covering the cathode. However. and the cathode article to e plated) is connected to the negative terminal. The cations are reduced at the cathode to deposit in the metallic. t the cathode. The u2+ associates with the anion S 42. These free cyanides facilitate anode corrosion. hen the e ternal power s pply is switched on. notably brass and solder. The plating is most commonly a single metallic element.. These cations associate with the anions in the solution. in an acid solution. For example. the metal at the anode is oxidized from the zero valence state to form cations with a positive charge. the u2+ is reduced to metallic copper y gaining two electrons. non-metal chemicals such as carbonates and phosphates may be added to increase conductivity.g. potassium cyanide) in addition to cyanides of the metal to be deposited. The anode is connected to the positive terminal of the s pply.a attery or. help to maintain a constant metal ion level and contribute to conductivity.     ! © ¦       ©¥ ¨¦¥ ¨§¦¥  !      ©    ¤ " £ ¢¡  . this means electrochemical oxidation of metals in reaction with an oxidant such as oxygen. the term degradation is more common Reaction It has been demonstrated that potential differences within a metal. Anodic reactions are typified by the dissolution of iron: Fe --> Fe+2 + 2eAnalogous reactions occur in other metals. commonly known as rusting. The electrons migrate through the metal to the cathode area where they react in any one of several ways. or between two metals. This type of damage typically produces oxide(s) and/or salt(s) of the original metal. In the most common use of the word. will cause chemical reactions at the anode and cathode. Corrosion can also refer to other materials than metals. such as ceramics or polymers. although in this context. Formation of an oxide of iron due to oxidation of the iron atoms in solid solution is a well-known example of electrochemical corrosion. .Corrosion Definition Corrosion is the disintegration of an engineered material into its constituent atoms due to chemical reactions with its surroundings. Positively charged ions will move toward the cathode.> 40H.----> MO Involves more noble metals in solution. Metal ion reduction (plating) +N + Ne. The most frequent cathodic reactions are a. Hydroxyl ions will combine with the ferrous cations produced by dissolution of the metal: Fe+2 + 2OH. c. aerated waters.> H2 Important in acidic solutions.> H2+2OH.--. d. Negatively charged ions. b.Important in natural.---.g. The floe is then rapidly oxidized to ferric hydroxide: 4Fe(OH)2 + O2 + 2H2O ----> 4Fe(OH)3 Dehydrolysis of this product leads to the formation of the corrosion products normally seen on ferrous surfaces.Occurs normally in natural waters.---.> 2H20 Occurs in aerated acidic solutions. Reduction of water 2H20 + 2e. red dust and hydrated ferric oxide: 2Fe(OH)3 ----> Fe203 + 3H2O Fe(OH)3 ----> FeOOH + H2O $ # .--. b. such as hydroxyl ions produced at the cathode. migrate to the anode of the corrosion cell.----> Fe(OH)2 The ferrous hydroxide produced has a very low solubility and is quickly precipitated as a white floe at the metal-water interface. Sulfate ion reduction 4H2 + S04-2 ---. c and d. turbulent conditions (e. acid cleaning).> S-2 + 4H2O Occurs in the presence of sulfate reducing bacteria.---. Ferric ion reduction Fe+3 + e. g.> Fe+2 Occurs under acidic. f. This movement of ions can cause additional reactions at the anode. Oxygen reduction 02 + 4H+ +4e. Oxygen reduction of water O2 + 2H2O + 4e. Hydrogen ion reduction 2 + + 2e.Some typical cathodic reactions are as follows: a. e. clay or microbiological slime. a tight. The degree to which such a film can impede corrosion is a complex function of the corrosion reactions. sand. silt. a corrosion film may show traces of hardness salts. Thus. or suspended matter like mud. because metal ions can penetrate it and reach the solution interface. the film will retard. or halt the corrosion. however. Initially. adherent film is formed. metal dissolution is not impeded by a film of corrosion products. corrosion can continue. ionic diffusion is prevented and the metal will no longer dissolve. the structure of the deposit and the water velocity . The structure of the entire surface film. they may cause the precipitation of other ions from the water. Most corrosion occurs at the beginning of a metal's service life. In time. If a porous film forms over the metal. including corrosion products and inclusions. If. is a major factor in determining the total amount of corrosion which will take place.As solid corrosion products are precipitated at the anode. discussing their general significance with respect to the mechanism of corrosion.Factors influencing corrosion Accordingly on this basis we list below some of the more important factors. Thick structural sections are more susceptible to corrosive attack that thin sections because variations in physical characteristics are greater. moisture-laden air is considerably more detrimental to an aircraft than it would be if all operations were conducted in a dry climate. the type. cause. Size and Type of Metal. Cause of corrosion. Climate. and seriousness of metal corrosion. Some of these factors can be controlled and some can not. In a predominately marine environment ( with exposure to sea water and salt air ). Factors Associated Mainly with the Metal y Effective electrode potential of a metal in a solution y Overvoltage of hydrogen on the metal y Chemical and physical homogeneity of the metal surface y Inherent ability to form an insoluble protective film Factors Which Vary Mainly with the Environment y Hydrogen-ion concentration (pH) in the solution y Influence of oxygen in solution adjacent to the metal y Specific nature and concentration of other ions in solution y Rate of flow of the solution in contact with the metal y Ability of environment to form a protective deposit on the metal y Temperature y Cyclic stress (corrosion fatigue) y Contact between dissimilar metals or other materials as affecting localized corrosion. The environmental conditions under which an aircraft is maintained and operated greatly affect corrosion characteristics. and postponing until later chapters the detailed discussion of others. When large pieces are machined or chemically milled after heat treatment. speed. . It is a less known fact that variations in size and shape of metal can indirectly affect is corrosion resistance. the thinner areas will have different physical characteristics than the thicker areas. many factors affecting corrosion. Temperature considerations are important because the speed of electrochemical attack is in creased in a hot. It is well known fact that some metals will corrode faster that others. moist climate. leaving the protected metal intact. Sn. by electroplating is also utilized to prevent corrosion. . The zinc also forms a protective coating of ZnCO3.. the exposed zinc again corrodes before the iron and continues to protect it. nickel etc. O2 and CO2. Galvanizing Coating iron or steel with a thin zinc layer is called 'galvanizing'. an alloy containing copper is another metal alloy which is less expensive and non reactive. Sacrificial protection Covering the surface with a more electro positive metal than Fe. Zinc preferentially corrodes or oxidizes to form a zinc oxide layer that does not flake off like iron oxide rust. Also. Dipping the iron/steel object in molten zinc and using it as the negative cathode zinc is coated on it. if the surface is scratched. Stainless steel is an example of a non-rusting alloy of iron and carbon. Alloying Iron or steel along with other metals can also be protected by 'alloying' or mixing with other metals (e. because the more reactive protecting metal is preferentially oxidized away. zinc or magnesium).Covering with zinc. Brass.Zn(OH)2. Prevention for iron rusting : Barrier protection The metal surface is not allowed to come in contact with moisture. chromium. chromium) to make non-rusting alloys. Example: Galvanization . The more electro positive metal loses electrons and as long as this coating is present Fe is protected. iii) Electroplating with non-corroding metals like Ni.g. This is referred to as sacrificial protection or sacrificial corrosion.g. ii) Applying oil or grease.Preventing corrosion Sacrificial Protection 'Rusting' can be prevented by connecting iron to a more reactive metal (e. Steel cans are protected by relatively un-reacted tin and works well as long as the thin tin layer is complete. Cr. i) Coating the metal surface with paint. Zn. Al. iv) Coating with alkaline phosphate (anti rust) solution.. This layer is produced by electrolytic deposition. Electroplating Coating the surface with metals like tin. Elva Kumalasari XI KI . The more reactive metal is the anode. Example: Fe can be connected to Mg. Used for protecting under ground pipes from rusting. It loses electrons and gradually disappears.Electrical protection (Cathodic protection) The iron object is connected to a more active metal either directly or through a wire. Zn or Al which are called the sacrificial anodes. Fe acts as the cathode.