L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.1 REDOX EQUILIBRIA II II. ELETROCHEMICAL CELLS電化電池 (A) Metal/Metal ion system — Qualitative study of Redox System When a metal is dipped into浸入 a solution containing ions of the same metal, the following three possible interactions may occur to different extents不同程度, depending on the nature of the particular metal/metal ion system: <i> <ii> A metal ion in solution may collide with the electrode but undergoes no change. A metal ion in solution may collide the electrode, gain electrons and can be deposited on the metal electrode. The metal ion is said to be reduced. The half equation is Mn+(aq) + ne M(s) and can be used to represent this reduction reaction. <iii> A metal atom from the electrode may lose electrons and enter the solution as the metal ion. The metal atom is said to be oxidized. The half equation is M(s) Mn+(aq) + ne and can be used to represent the oxidation reaction. An equilibrium is established between the tendencies of <ii> and <iii>. Mn+(aq) + ne M(s) In the forward process, that is the reduction process, the hydrated metal ions in solution consume electrons from the electrode. This results in a net deficit缺少of electrons in the electrode and thus a positive charge on the electrode. On the other hand, in the reverse process, that is the. oxidation process, atoms from the metal lattice pass into solution to form cations. This leaves a surplus過剩of electrons and results in a negative charge on the electrode. The charge on the electrode depends on which process occurs more readily. If <a> the forward process is more favourable, reduction predominates and thus the electrode acquires a positive charge. <b> the reverse process is more favourable, oxidation predominates and thus the electrode acquires a negative charge. In either case, there is a separation of charge and thus a potential difference is set up between the metal and its ions in the solution. The metal/metal ion system is known as a half cell. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.2 (B) Use of Metal/Metal Ion System as Sources of Electrical Energy The combination of two metal/metal ion half-cells will give an electrochemical or voltaic cell. An electrochemical cell is a device which produces an electromotive force (e.m.f.) 電動勢as a result of chemical reactions taking place at the electrode. An electrochemical cell is thus a device which convert chemical energy into electrical energy. Each cell consists of two half cells. In one half-cell an oxidation half-reaction takes place. In the other a reduction half-reaction takes place. Example : The Daniell cell丹聶耳電池 Salt bridge Zn Cu Cotton wool Daniell cell In the Daniell cell, zinc and copper electrodes are dipped into a 1M solution of zinc sulphate and copper(II) sulphate solution. The function of the porous partition of the salt bridge is to enable electrical contact while separating the two different electrolytes, so that direct chemical redox reaction Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) will not occur. Electrical contact between the solutions is established in the walls of the porous partition, or through the electrolytes in the salt bridge. Salt bridge : It can be made by an inverted U-tube containing saturated salt solution such as potassium chloride, potassium nitrate or ammonium nitrate. I When these two half cells are connected externally by a conducting wire, owing to the difference in potentials between the half cells, an electrochemical reaction involving the passage of electrons can take place. <1> The zinc goes into the solution as zinc ions. Oxidation takes place. _____________________________________________________ <2> The electrons produced flow in the external circuit to the copper electrode. In the copper electrode, copper(II) ions in solution accept these electrons and are deposited. _____________________________________________________ L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.3 The zinc electrode become the negative terminal of the cell whereas the copper electrode becomes the positive terminal of the call. The Daniell cell therefore provides a means of converting the chemical potential energy between the two metal/metal ion system into electrical energy. The enthalpy of the net reaction Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) is liberated 釋放 in the form of electrical energy as electron flow. (C) Potential difference of electrochemical cells When a metal is placed in contact with a solution of one of its salts, an equilibrium is established between the tendencies. oxidation M(s) Mn+(aq) + ne reduction The potentials set up will differ for different metals. These potentials (absolute potentials) cannot be measured, but the difference in potential between two metal/metal ion system can be found by incorporating them into an electrochemical or voltaic cell, and measuring the potential difference between the metal electrodes. Note <1> As the resistance in the external circuit is increased, the potential difference between the two systems increases accordingly. whereas the current flowing through (measured by the milliammeter) is decreases. The Maximum potential difference of the whole cell is known as the electromotive force電 動勢 or e.m.f. and it occurs when no current is flowing. The value of this e.m.f. is a measure of the relative tendencies of the electrode system involved to liberate electrons by forming ions in solution. <2> Such set up can be used to measure the e.m.f. values of different metal-metal ion systems, and their value compared to assess the relative tendencies of metal/metal ion system to release electrons. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II (D) Chpt:20 p.4 Cell Diagrams 電池圖 A cell digram can be used to represent the electrochemical cell and its e.m.f.. Example : The Daniell cell can be represented as follows: Zn(s) Zn2+ (aq) Cu2+ (aq) Cu(s) E = +1.1V <1> The solid vertical lines represent boundaries between the electrode and its solution in each electrode system. <2> The dotted lines represent the salt bridge or porous pot. Usually, represents a salt bridge whereas represents a porous <3> E represent the e.m.f. of the cell in volts. The (+) or (-) sign indicates the polarity of the right hand electrode. In the above case, copper electrode is the positive pole of the cell. <4> By convention, the half-cell with the cathode陰極 (the positive electrode) is placed on the right. This is the half-cell in which reduction takes place. Thus, Cu2+(aq) Cu(s) represents Cu2+(ag) + 2e Cu(s) The half-cell with the anode 陽極 (the negative electrode) is placed on the left. This is the half-cell in which oxidation takes place. Thus, Zn(s) Zn2+(aq) represents Zn(s) Zn2+(aq) +2e (E) Types of half-cells The three most common types are metal-metal ion, non-metal-ion and ion-ion half cells. 1. The Metal-Metal Ion Half-Cell The zinc-zinc ion and copper-copper(II) ion electrodes. 2. The Non-Metal-Ion Half-Cell The iodine-iodide ion electrodes (with graphite as the inert electrode). The following equilibrium is established. I2(aq) + 2e 2I-(aq) Cell diagram : <a> As a cathode <b> As an anode 3. The Ion-Ion Half-Cell The iron(III)-iron(II) half cell. In this the following equilibrium is established. Fe3+ (aq) + e Cell diagram : <a> As a cathode Fe2+ (aq) <b> As an anode L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II III. ELECTRODE (A) Chpt:20 p.5 POTENTIALS 電極電勢 Standard Electrode Potential標準電極電勢 If the metal is immersed in a solution of its ions of concentration 1 mol dm3 at 25°C, then the potential acquired under these standard conditions if the standard electrode potential of the metal. E0. Examples : E0Cu = +0.34V Note: (B) E0Zn = -0.76V Since the electrode potential depend on temperature, concentration and also pressure, it is necessary to standardize them if they are to be compared. Finding Electrode Potentials If a cell is constructed from a standard electrode (i.e., one of known potential) and an electrode of unknown potential, the e.m.f. of the cell can be used to find the unknown electrode potential. For a cell with the standard electrode on the left- hand side E0cell = Eelectrode of unknown potential - Estandard electrode (C) Standard Hydrogen Electrode 標準氫電極 The standard hydrogen electrode is the reference electrode with which other electrodes are compared. It consists of a platinised platinum electrode immersed in a solution of 1 mol dm-3 hydrogen ions. Hydrogen gas at a pressure of 1 atm is bubbled over the platinum electrode. On the surface of the platinum, equilibrium is established between hydrogen gas and hydrogen ions. 2H3O+(aq) + 2e 2H+(aq) + 2e 2H2O(l) + H2(g) or simply H2(g) The potential is assigned a value of 0 volts. In a cell diagram, a hydrogen electrode is represented as ____________________________ if platinum is the anode ____________________________ if platinum is the cathode The standard electrode potentials of other system can be found by combining these with a standard hydrogen electrode and measuring the e.m.f. of the cell formed. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Example : Chpt:20 p.6 A voltaic cell which combines a standard zinc electrode and a standard hydrogen electrode. The two compartments are connected by a salt bridge The e.m.f. of the cell is -0.76V. The voltmeter shows that electron flow through the external circuit from zinc to the hydrogen electrode, showing that zinc has a standard electrode potential of -0.76V That is Ecell = -0.76V, Estandard electrode = 0V Eelectrode of unknown potential = Ecell + Estandard electrode = (-0.76V) + 0 = -0.76V (D) The e.m.f. of electrochemical cells As the electrochemical cell consists of two half-cells. Each half-cell has its own electrode potential. When these two half-cells are connected into an external circuit, two important things happen. 1. The cell redox reaction begins to occur. This reaction is the sum of the two half-reaction. As a result, the equilibrium in the two half-cells is disturbed. As the cell redox reaction proceeds, the concentration of the oxidized species in the half-cell containing the anode increases. On the other hand, the concentration of the oxidized species in the half-cell containing the cathode decreases. Eventually the redox reaction reaches equilibrium, when this occurs the battery has run down or is “flat”. 2. When the two half-cells are connected into an external circuit, current flows from the negative electrode, that is the anode, to the positive electrode, that is the cathode. The current is driven by the potential difference. In the external circuit the current is a flow of electrons. In the two half- cells the current is carried by the ions. As current is taken from the cell, the concentration of the oxidized species increases and the concentration of reduced species decreases. At equilibrium, the potential difference between the two electrodes is zero. The maximum value of potential difference between the two electrodes is known as the electromotive force or e.m.f. of the cell. E0cell = E0anode – E0cathode L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II (E) Chpt:20 p.7 Measuring Cell e.m.f. and Electrode potentials Accurate measurement of the e.m.f. of a cell by a voltmeter is impossible since this instrument draws current from the cell. This causes electrode reactions to occur, thereby changing the concentration of the solution and altering the potential difference. The problem can be overcome by the use of a high resistance voltmeter, or just by a digital voltmeter so that the current taken is small. By knowing the e.m.f. of the cell, the electrode potential of a metal/metal ion system (a half cell) can be determined by setting up an electrochemical cell with the hydrogen electrode as the other half cell. If the hydrogen electrode is the negative electrode, the cell diagram is Pt H2(g) H+ (aq) Mn+ (aq) M(s) Thus E0cell = E0 (Mn+/M) – E0( H+/H2) = E0 (Mn+/M) Note: The standard electrode potential values measured may have positive or negative values in volts. A positive value implies that when this electrode is coupled with the standard hydrogen electrode, the original electrode will be the cathode and reduction occurs. A negative value implies that when this electrode is coupled with the standard hydrogen electrode, the original electrode will be the anode and oxidation occurs. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.8 Exercise 1 From the following data, determine (a) the standard cell e.m.f. (b) the cell reaction (c) the cell diagram. Data : (F) Zn2+(aq) + 2e Zn(s) E0(Zn2+/Zn) = -0.76V Ni2+(aq) + 2e Ni(s) E0(Ni2+/Ni) = -0.26V Standard Electrode Potentials: Extended Redox Potential Series Electrode potentials of half-cells are often called redox potentials. By convention, redox potentials are always quoted as reduction potentials還原電勢. The equilibrium half-reactions are thus written reduction as the forward reaction: Oxidized species + ze Reduced species A reduction potential is a measure of the tendency for reduction to occur. Besides involving a metal/metal ion for redox system, there are other redox systems involving reactions between non-meals and non-metals, and also between ions only. Examples : Br2(aq) + 2e 2Br-(aq) Fe3+ (aq) + e Fe2+ (aq) E0 = +1.065V E0 = +0.77V Note : <1> An inert platinum electrode is often inserted into these half-cells system to provide a pathway for electrical conduction. <2> When writing the half-cell diagram, the reduced form of the ion is put near the inert electrode, and separated it from the oxidized form by a comma. Br2(aq),,2Br-(aq)Pt or Pt 2Br-(aq),Br2(aq) Fe3+(aq) ,Fe2+(aq) Pt L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.9 <3> When the oxidized and reduced forms of an electrode system contain more than one chemical species (ion or molecule) which participate in the cell reaction, These ions or molecules should be included in the oxidized and reduced forms of the half cell. Example: MnO4-(aq) + 8H+(aq) + 5e Mn2+(aq) + 4H2O(l) E0 = +1.51V The cell diagram should be [MnO4- (aq)+8H+(aq)],[Mn2+(aq) + 4H2O(l) ] Pt(s) Extended Reduction Potentials The electrode potentials of the various systems, when tabulated together with those of the metal/metal ion systems, will form the extended reduction potential series. Interpretation <1> In such series, electrode system with the greatest negative electrode potential are at the top of the list, and they have the greatest tendency to exist as cations in solution. That is, they are strongest reducing agent. <2> Such electrochemical series shows the order of decreasing reactivity of the system, an also shows which elements can displace each other from solutions of their salts. (G) The electrochemical series – Reduction potential series When elements are placed in order of their standard electrode potentials , the electrochemical series is obtained. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.10 <1> The series can be directly obtained from the redox potential series. That part of the series which includes metals only is called the electrochemical series of metals. This series closely approximates to the order of reactivity of metals. <2> The most electropositive and most reactive metals are at the top of the series. These metals are readily oxidized and thus ionize easily. As a consequence they are strong reducing agents.. Metals at the bottom of the series do not oxidize readily. They are thus stable in their reduced form. Gold and Mercury are examples. <3> DISPLACEMENT REACTIONS置換反應 Metals high in the series reduce the oxidized forms of metals lower in the series. In effect they displace the metals from their oxides or from solutions of their salts. Example 1: Zinc displaces copper from a solution of copper(II) sulphate. Zn(s) + Cu2+ (aq) Zn2+ (aq) + Cu(s) On the other hand, magnesium, which is higher in the series, displace zinc from its oxides: Mg(s) +ZnO(s) MgO(s) + Zn(s) Example 2: Metals above hydrogen in the electrochemical series of elements reduce hydrogen ions to form hydrogen gas. Mg2+(s) + 2H+ (aq) Mg (aq) + H2(g) Example 3: While the reactivity of metals increases as the electrode potentials become more negative, the reactivity of non- metals increases as the electrode potentials become more positive. Chlorine displaces iodine from solution. Cl2(g) + 2I-(aq) 2Cl-(aq) + I2(s) (H) Uses of Standard Reduction potentials Values of standard reduction potentials can be used to predict possible redox reactions預測氧化 還原反應的可行性 and to calculate the e.m.f. of cells. A redox reaction can occur spontaneouly自發地進行 (feasible to occur可行) when the e.m.f. calculated is a positive va1ue. Note The value of the e.m.f. is only a state function. It only shows that the reaction is energetically feasible or not by inspecting the sign of the value. It could not tell the rate of the reaction since the rate depends on the kinetic factors such as the activation energy of the reaction. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II (I) Chpt:20 p.11 Feasibility of redox reaction 氧化還原反應的可行性 Thermodynamics predicts that certain reactions are capable of occurring whereas others are not. Reactions capable of occurring are said to be feasible. However, although a reaction may be feasible it may not necessarily occur spontaneously. Most combustion reactions do not occur spontaneously although they are thermodynamically feasible. This is because the energy barrier has to be overcome before the reaction can start. So, under given conditions a reaction may be thermodynamically feasible but not kinetically feasible. The thermodynamic feasibility of a redox reaction can quite simply be determined by inspecting the electrode potentials of the two half-reactions. A reduction half-reaction is feasible if its electrode potential is more positive than the electrode potential of the other half- reaction. Reduction is not feasible if the electrode potential is less positive than that of the other half-reaction. The half-reaction with the less positive electrode potential must be oxidation half-reaction. Exercise 2 Given that : Al3+(aq) + 3e Al(s) E° = -1.66V Cu2+(aq) + 2e Cu(s) Eo = -0.34V (a) Will aluminium metals displace copper(II) ions from the aqueous solutions? (b) Can we tell the rate of the displacement reaction from the values of electrode potentials? ANSWER . Exercise 3 (a) Given that Br2(aq) + 2e 2BrO2 (g) + 2H2O(l) + 4e E0= +1.09V 4OH-(aq) E0= +0.40V Predict whether Br2 or O2 is a stronger oxidizing agent? How can such prediction be demonstrated experimentally? (b) Given that Ag+(aq) + e Ag(s) Eo= +0.779V AgC1(s) + e Ag(s) + Cl-(aq) Eo= +0.220V O2(g) + 2H2O(l) + 4e 4OH-(aq) Eo= +0.401V Show that silver metal will not undergo oxidation in the presence of oxygen and water only, but will do so when chloride ions are present. ANSWER L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.12 Exercise 4 Given that Ag (aq) + e Ag(s) E = +0.80V AgC1(s) + e Ag(s) + Cl (aq) E°= +0.22V Devise an electrochemical cell in which a spontaneous reaction can take place. Write the cell diagram for this cell. ANSWER IV. COMMERCIAL CELLS 商用電池 A battery is portable source of electricity. It consists of one or more electrochemical cells. A cell which does not regenerate reactants is called primary cell 原電池. (e.g. dry cell) A cell which can be recharged by regenerating the cell reactants is called secondary cell 次電池 or storage cell. (e.g. lead acid accumulator鉛酸蓄電池) In each type of cells, redox reaction should occur when it generates electricity. (A) The Dry cell The dry cell is a primary cell which is the commonest, cheapest and most convenient cell used at present. All its components are solids or pastes which are tightly sealed from the environment. The anode is the zinc container which encases the dry cell. The cathode is a graphite rod surrounded by a layer of manganese(IV) oxide and carbon. The electrolyte is a paste consisting of zinc chloride, ammonium chloride and water. The half-reactions at the electrode are at the anode at the cathode The overall equation is Zn(s) + 2MnO2(s)+ 2NH4+(aq) Zn2+(aq)+ Mn2O3(s) + 2NH3(g)+ H2O(l) Note : <1> <2> A dry cell generate between 1.25V and 1.5V. The formation of ammonia around the cathode would disrupt the electric current. This is prevented by the complex formation with Zn2+ 4NH3 + Zn2+ Zn(NH3)42+ L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II (B) Chpt:20 p.13 Secondary cell- Lead storage cell (Lead-acid accumulator) Lead accumulator is a secondary cell, the half-reactions at the electrodes are readily reversible. It consists of a lead anode arid a grid of lead packed with lead(IV) oxide as the cathode. The electrolyte is sulphuric acid. The half-reactions at the electrodes are at the anode: at the cathode The overall reaction is Pb(s)+ PbO2(s)+ 4H+(aq)+ 2SO42-(aq) discharging 2PbSO4(s)+ 2H20(l) charging Note <1> One cell provides about 2V. The storage battery used in cars normally consists of 6 these cells arranged in series to provide 12V. <2> The battery is recharged by applying a current from an external source. This reverses the electrode reactions. The state of charge can be indicated by the relative density of sulphuric acid (i) When the storage cell is fully charged, the sulphuric acid has a relative density of 1.275. (ii) When the cell is used for some time, the concentration of sulphuric acid will decrease and also its relative density. Upon discharging, lead (II) sulphate forms at both electrodes, thus reducing the concentration and relative density of the sulphuric acid. The relative density thus indicates the state of charge in the battery. Lead acid accumulator L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II (C) Chpt:20 p.14 Fuel cell 燃料電池– Hydrogen-oxygen fuel cell A fuel cell is a type of primary cell in which the reactants are continuously replaced as they are consumed and the products are continuously removed. In the hydrogen-oxygen fuel cell, hydrogen and oxygen are bubbled through porous carbon electrodes into a concentrated solution of an alkali. The carbon electrodes contain a platinum catalyst. The half-reactions at the electrodes are at the anode 2H2(g) + 4OH-(aq) 4H2O(l) + 4e at the cathode O2(g) + 2H2O(l) + 4e 4OH- (aq) The overall reaction is 2H2(g) + O2(g) 2H2O(l) The water is removed and in spacecraft is consumed by the astronauts as the drinking water. Note: In some fuel cells an acidic electrolyte is used in which case the electrode reactions are at the anode : H2(g) 2H+(aq) +2e at the cathode : O2(g) + 4H+(aq) + 4e 2H2O(l) The overall reaction is 2H2(g) + O2(g) 2H2O(l) L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II V. Chpt:20 p.15 CORROSION OF IRON AND ITS PREVENTION鐵的腐蝕與防預 (A) Introduction Corrosion may be regarded as the natural tendency of metals to return to their oxidized species. Corrosion is a redox process. reduced species (metal) oxidized species + ne The main sources or causes of corrosion are <1> the atmosphere - humidity and atmospheric pollution, (most important) <2> submersing 浸入 in water of aqueous solution, <3> underground soil attack, <4> corrosive gases, <5> immersion in chemicals. Corrosion of iron is called rusting. Rusting is the electrochemical process by which iron rust away when exposed to air and water. (B) The electrochemical process involved in rusting Rusting is the most common example of corrosion. When iron comes into contact with acid, oxygen an other species found in the environment, electrochemical reactions takes place. <1> Consider an iron sheet exposed to open atmosphere. The plate may have other metal impurities incorporated in it When the plate gets wet due to moisture in the air, the thin layer of water tends to dissolve oxygen, which diffuses from the edge of the water drop to the interior. <2> Around the edge of the drop of water 水滴的邊沿, the concentration of atmospheric oxygen is higher, and the following reaction may occur on the iron surface immediately in contact with the edge of the water drop. 1/2 O2(g) + H2O(l) + 2e 2OH-(aq) cathodic reaction Note 1. If another metal impurities (which accepts electrons more readily than iron) happens to be present at this site, then the above reaction may occur even more readily. 2. The reaction is favored by a low pH, so that the presence of acid , or even carbon dioxide, will accelerate rusing. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.16 <3> At the centre of the water drop, the concentration of atmospheric oxygen is lower, so that the following reaction takes place preferentially. Fe(s) Fe2+(aq) + 2e anodic reaction Note: 1. Since this process involved release of electrons, it takes more readily on non-uniform surface (e.g. pointed ends) and mechanically strained areas, where the electrons are of higher energy and can be released more readily. 2. The presence of a more electropositive metal as impurities also favors this process, as it releases electrons even more readily than iron. <4> The free Fe2+(aq) and OH-(aq) diffuse from their sites of formation and precipitates as Fe(OH)2 precipitate, which is further oxidized by dissolve oxygen to form iron(III) hydroxide. On standing, this changes to rust (Fe2O3.nH2O), a reddish brown solid. In many cases this coating either flakes off or is permeable to both water and air, so that corrosion continues unhindered. NOTE: (1) The whole process is accelerated by electrolytes like sodium chloride, which gently increases the conductivity of the solution. e.g. Iron window frames by the seashores rust more readily as the thin layer of water on iron surface contains dissolved sodium chloride from salts spray near the sea. (2) The whole process is also accelerated by high temperature, which increases the rate of a chemical reaction. e.g. Car exhaust pipes rust easily as they are subjected to great heat. They have to be replaced every couple of years. SUMMARY Corrosion of iron (rusting) is a local electrochemical process, with the metal at the centre of the water drop being the anodic site, whereas the metal (or other metal impurties in contact with it) at the periphery of the drop being the cathodic site. The corrosion process of iron can be proved by considering their electrode potentials. Fe2+(aq) Fe(s) E0 = -0.44V anode 0 Pt O2(g) OH (aq) E = - 0.4V cathode E0cell = (-0.40V) - (-0.44V) = 0.04V As the value of E0 is positive, the above cell is spontaneous. The whole process can be represented by a cell diagram Fe(s) Fe2+(aq) OH-(aq) O2(g) metal impurities(s) or Fe(s) Such a local voltaic cell is formed by iron together with other metal impurities, as well as water and air. Dissolved electrolytes, e.g. NaCl and HCO3- from atmosphere, a low pH and a high temperature, will accelerate the process. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.17 (C) Prevention of Corrosion Various methods are used to protect a metal from corrosion. <1> Coating. Coating the metal with paint or tin can prevent attack by atmospheric moisture and oxygen. However, if coating is flawed or broken, corrosion occurs. e.g. Plating with a thin layer of tin. When a tin-coated iron can containing lemon juice is scratched. The two metals are exposed to the acid, an electrochemical reaction would occur. Fe2+/Fe = -0.44V 2+ Sn /Sn = -0.136V Iron is more reactive than tin so it tends to reduce any tin ions that are present. Thus, the tin ions aids the oxidation of iron. As a result, the scratched tin-plated object has a stronger tendency to corrode. <2> Cathodic protection The use of zinc to protect iron form corrosion is called cathodic protection. When iron is coated with zinc, it is called galvanized iron. A galvanized iron can containing acidic rain water is scratched. The iron surface is exposed. Some metal ions form in solution. Since E0 : Fe2+/Fe -0.44V Zn2+/Zn -0.76V Zinc is more electropositive than iron (more likely to form cation) than iron so that it tends to reduce any iron ions that form. Zn(s) + Fe2+(aq) Zn2+(aq) + Fe(s) So the presence of zinc prevents the corrosion of iron. Iron will be protected from rusting even the galvanized iron 鍍鋅鐵 is scratched 刮花. Note: <1> Since Zn is higher in the electrochemical series than iron, it acts as the anode and the iron acts as the cathode. The zinc is thus oxidized in preference to the iron. Zn(s) Zn2+(aq) + 2e <2> Zinc is not used in food canning since it is higher than tin in the electrochemical series. It is thus more susceptible to attack by acids in fruit juice. Besides, Zn2+ ion is toxic. <3> A sacrificial anode is also used as a means of cathodic protection. A metal (e.g. magnesium ) much higher than iron in the electrochemical series is chosen as the sacrificial metal. Such metal has a higher tendency to release electron than iron, it always form anode whenever an electrochemical cell is set up between it and iron. L.S.T. Leung Chik Wai Memorial School F.6 Chemistry Chapter 20: Redox Equilibria II Chpt:20 p.18 Mg(s) Mg2+(aq) + 2e (oxidation) The sacrificial metal forms the anode whereas the iron surface forms the cathode to accept the electrons released from anode. In practice, the sacrificial metal is connected to the object (e.g. steel pipe) through a conducting wire. Iron is protected from corrosion as the metal remains. <4> Aluminium, although high in the electrochemical series, does not corrode readily since it oxidizes to form aluminium oxide, A12O3, which does not flake. It thus forms a strong protective coating. (C) Socioeconomic implication of corrosion and prevention As a construction material, iron is suffered from the drawback that it corrodes quite readily. Corrosion of iron can cost much to the society. In terms of money, it is estimated in 1980 in the US that 70 billion dollars were lost annually because of corrosion. Apart from money spent on the replacement of the corroded articles and the prevention methods, there are indirect costs such as those for the maintenance of machines, and those due to lost production when machines fail down or when they are shutdown for maintenance. The wastage of natural resources is one of the social implications of corrosion. It has been estimated that 1 tonne of steel is converted into rust every 90 seconds in Britain, and that about 40% of the steel made in the US is used to replace steel lost by rusting. Other than wasting natural resources, corrosion of iron causes considerable inconvenience to human beings, and even lost of life. This happens because corrosion results in the formation of cracks and crevices which weakens the strength of metals. Concrete in buildings may fall off if the steel reinforcing bars inside the concrete corrode. Moreover, there is evidence that there may be a correlation between the number of serious injuries suffered in road accidents and the age (i.e. amount of corrosion) of the vehicle. Despite of the disadvantages, iron is still the most important in present-day society. The disadvantages of corrosion are outweighed by the relative low cost, abundance and ease of extraction of iron. Replacement of iron by a more corrosion-resistant metal of a reasonable price will be the goal of many metallurgists. But, more than they should be done, the society as a whole should assume a greater responsibility for corrosion prevention in order to have a greater improvement of the situation.