UNIT-II ELECTROCHEMISTRY AND BATTERIES. Electrolysis: The changes in which electrical energy causes chemical reaction to occur.

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1 ELECTROCHEMISTRY AND BATTERIES Concept of electrochemistry: The branch of science which deals with the relationship between electrical energy and chemical energy and their inter-conversion of one form to another is called electrochemistry. Electrolysis: The changes in which electrical energy causes chemical reaction to occur. Electro chemical cells: The changes in which electrical energy is produced as a result of chemical change. Conductors: The substances which allow the passage of electric current are called conductors. Eg. Metals like Cu, Ag, Sn etc. Insulators: The substances which do not allow the electric current to pass through them are called non-conductors or insulators. Eg. Rubber, wood, wax, wool, glass etc. Metallic conductors: These are the metallic substances which allow the electric current to pass through them without undergoing any chemical change. The flow of electric current in metallic conductors is due the flow of electrons in the metal atoms. Eg. Metals like Cu, Ag, Sn etc. Electrolytes: These are the substances which allow the electric current to pass through them in their molten states or in the form of their aqueous solutions and undergo chemical decomposition. The flow of electric current through an electrolytic solution is called electrolytic conduction in which charge is carried by ions. These substances do not conduct electricity in the solid state but conduct electricity in the molten state or aqueous solutions due to movement of ions. Eg. Acids, bases and salts. Non-electrolytes: The substances which do not conduct electricity either in their molten state or through their aqueous solutions are called non-electrolytes. Eg. Sugar, glucose, urea, ethyl alcohol etc. Strong electrolytes: The electrolytes which are almost completely dissociated into ions in are called strong electrolytes. Eg: NaCl, KCl, HCl, NaOH, NH4NO3 Weak electrolytes: The electrolytes which do not ionize completely in solution are called weak electrolytes. Equilibrium is established between unionized electrolyte and the ions formed in the solution. Types of Electrodes Electrode potential (or) Single Electrode potential (E): It is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt. Standard Electrode potential (E ): It is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt of one molar concentration at 25 C. Types of Electrodes: An electrochemical cell consists of two electrodes, positive and negative. Each electrode constitutes a half cell or a single electrode. 1

2 CONSTRUCTION AND FUNCTIONING OF SOME STANDARD ELECTRODES: The single electrode potential is conveniently measured by combining the half cell with a standard electrode and measuring the total EMF of the cell. Ecell = Eright - Eleft Ecell = E right - E o (if standard electrode is anode) Ecell = E o - Eleft (if standard electrode is cathode) Where E o is the standard electrode potential. Eg. Standard hydrogen electrode, Calomel electrode, Quinhydrone electrode 1. Standard Hydrogen Electrode (S.H.E) Standard Hydrogen Electrode (S.H.E) is a primary reference electrode. Construction: S.H.E consists of a platinum electrode coated with Pt-black immersed in 1 Molar solution of H + ions maintained at 25 o C. Hydrogen gas at 1atm pressure is constantly bubbled into the glass hood over the Pt-electrode. Hydrogen gets oxidized at the Pt-electrode and passes into the solution forming H + ions and electrons. The standard electrode potential of Hydrogen electrode at these conditions is considered as zero. Depending upon the half cell to which it is combined, hydrogen electrode can act as either anode (or) cathode. When S.H.E acts as anode: H2(g) 2H + + 2e - When S.H.E acts as cathode: 2H + + 2e - H2(g) Therefore S.H.E is regarded as reversible electrode. H2(g) 2H + + 2e - Applications of S.H.E: Determination of ph for the unknown solution: For a Hydrogen electrode at 298K, the Nernst equation is given as follows: 2H + (aq) + 2e - H2(g) H + (aq) + e RT [H ] E = E - log nf [H ] 1/2 o 2 + H2(g) 1/2 o E = - log ( [H + 2] = 1 and E = 0) + E = (-log[h ]) nf [H ] E = ph 2

3 In order to determine the ph of unknown solution, a Pt-electrode is immersed into the solution and connected to S.H.E. representation: Pt, H2(g) H + (aq) (?) H + (aq) (1M) H2(g) (1atm), Pt E = E - E cathode anode E = E - E E E S.H.E H + 2 /H = 0- ( pH) = pH E ph = Calomel Electrode Calomel Electrode is a secondary reference electrode. Construction: It is a mercury-mercurous chloride (Hg-Hg2Cl2) electrode. The calomel electrode is used as only reducing electrode i.e. as cathode only. It consists of a glass tube having a side tube on each side. One side tube acts as salt bridge and other is used to fill KCl solution. The high purity mercury is placed at the tip of this tube and connected to the circuit by the means of Pt-wire, sealed in a glass tube. The surface of Hg is covered with a paste of mercurous chloride (calomel) and Hg in KCl solution. The electrolyte is the solution of KCl. The electrode is connected with the help of side tube on the left through salt bridge with the other electrode whose potential has to be determined. The potential of the calomel electrode depends upon the concentration of KCl. Representation of calomel electrode: Hg, Hg2Cl2(s) KCl (saturated solution) Concentration of KCl 0.1N N Saturated Electrode potential (V) 3

4 Determination of ph of unknown solution by Saturated Calomel electrode: representation: Pt, H2 (1atm) H + (?) KCl (saturated) Hg2Cl2, Hg E = E - E E E Calomel H + 2 /H = ( pH) = pH E ph = Ion Selective Electrode- Glass electrode They possess an ability to respond to only certain specific ions thereby developing a potential with respect to that species only in a given mixture and ignore all the other ions totally. The potential developed by ion selective electrode depends on the concentration of specific ion of interest. Eg: Glass electrode specific for H + ions and pressed pallet of Ag2S + AgCl specific for Cl - ions. Glass electrode: Construction: When two solutions of different ph values are separated by a thin glass membrane, a potential difference develops between the two surfaces of the membrane. This potential difference developed is proportional to the difference in ph value. The glass membrane functions as ionexchange resin.an equilibrium is established between Na + ions of the glass and H + ions in the solution. The potential of the glass electrode is given by: EG = E o G ph Where ph range of the test solution is The glass electrode consists of a thin walled glass tube containing coated Ag-electrode or simply Pt-electrode in 0.1M HCl. AgCl representation of glass electrode is given by: Ag AgCl(s) H + (0.1M) Glass or Pt, HCl (0.1M) Glass 4

5 Glass electrode & determination of ph by glass electrode: HCl in the bulb furnishes constant H + ion concentration. Thus it is Ag-AgCl electrode reversible with respect to Cl - ions. It is used as internal reference electrode for determining the ph of the solution especially the colored solution containing oxidizing or reducing agents. representation: Pt, HCl (0.1M) Glass KCl (sat) Hg2Cl2(s) Hg E = E - E o o Calomel Glass E = [E pH] E = o Glass E - E o Glass Since the resistance is very high, a special electron tube voltmeter is used to measure the EMF of the above cell. Advantages: 1. It is simple and can be easily used. 2. The results are accurate. 3. Equilibrium is rapidly reached. Limitations: 1. The range of solution ph is 1 to The resistance of the membrane is very high. Nernst Equation Standard electrode potentials are measured at standard states i.e at 1M concentration of electrolyte, temperature at 298 K and 1atm pressure. However the electrode potentials depend upon concentration of electrolyte solutions and temperature. Nernst equation gives the relationship between electrode potentials and concentration of electrolytic solutions. For the general reduction reaction occurring at an electrode, M n+ (aq) + ne - M(s) E = E 2.303RT [M ] - log nf [M ] n+ o n+ (s) M /M M /M n+ (aq) 5

6 Where E = Standard EMF of the cell for 1 M solution at 298 K; R = Gas constant; T = Kelvin temperature; n = number of electrons involved in the cell reaction; F = Faraday of electricity; E = electrode potential of the metal; [M] = activity of metal in the metal phase and is taken as unity; [M n+ ] = activity of metal ions in the solution is taken equal to their molarities; But at STP conditions, R = 8.314JK -1 mol -1 ; T = 298K; F = coulomb charge. Then the Nernst equation becomes o x x n+ E n+ n+ M /M = E M /M - log considering [M n+ (s)] = 1 n x [M ] E = E - log n [M ] n+ o M /M M n+ /M n+ (aq) o n+ E n+ n+ M /M = E M /M + log[m ] n Nernst equation enables us to calculate the following: 1. Half cell potential or single electrode potential. 2. potential or EMF of the cell. 3. Equilibrium constant for the cell reaction. Electrochemical series and its applications Definition: When the various electrodes (metals) are arranged in the order of their increasing values of standard reduction potential on the hydrogen scale, then the arrangement is called electrochemical series. (or) When the various electrodes (metals) are arranged in the order of their decreasing values of standard oxidation potential on the hydrogen scale, then the arrangement is called electrochemical series. (aq) Applications of electrochemical series: 1. Calculation of standard EMF of the cell. 2. Relative ease of oxidation (or) Reduction. 3. Displacement of one element by the other. 4. Hydrogen displacement behavior. 5. Corrosion tendency of the elements. 6

7 Electro chemical cell An electrochemical cell is a device which converts chemical energy into electrical energy. The redox reaction is utilized for generation of electrical energy. The electrochemical cells are commonly referred as Voltaic or Galvanic cells. The electrochemical cell is divided into two half cells. The half cell electrode where oxidation (loss of electrons) occurs is called anode (negative electrode). The half cell electrode where reduction (gain of electrons) occurs is called cathode (positive electrode). An electrochemical cell is the coupling of these two half cells. Construction of Daniel cell / voltaic cell The Daniel cell is a typical example of Galvanic cell. Daniel cell consists of a beaker containing copper rod dipped in CuSO4 solution which is connected to another beaker containing zinc rod dipped in ZnSO4 solution by a salt bridge. Salt bridge is an inverted U-tube containing saturated solution of some electrolyte such as KCl, KNO3, and NH4NO3 which does not undergo chemical change during the process. The saturated solution is generally taken in agar-agar jelly or gelatin. 7

8 The salt bridge allows the flow of ions to pass through it when the flow of electric current takes place. It completes the electrical circuit and maintains the electrical neutrality of two half cell solutions. When the circuit is completed, electric current flows through the external circuit as indicated by an ammeter. The following observations are made: 1. Zinc rod gradually loses its weight. 2. The concentration of Zn +2 (aq) in the ZnSO4 (aq) solution increases. 3. Copper gets deposited on the electrode. 4. The concentration of Cu +2 (aq) in CuSO4 (aq) solution decreases. 5. The flow of electrons is from Zn-electrode (anode) to Cu-electrode (cathode). 6. The flow of electric current is from Cu-electrode (cathode) to Zn-electrode (anode). The above observations can be explained as follows: Zn is oxidized to Zn +2 ions which go in the solution and therefore Zn rod gradually loses its (oxidation half cell). weight Zn Zn e - (At anode) The electrons released at the Zn electrode move towards Cu electrode through external circuit. These electrons are accepted by the Cu +2 ions in the solution and get reduced to copper which gets deposited on the Cu-electrode (reduction half cell). Cu e - Cu (At cathode) The oxidation of Zn occurs at anode (negative terminal) and reduction of Cu +2 occurs at cathode (positive terminal). The flow of electrons is from negative terminal (anode) to positive terminal (cathode). Flow of electrons (-) Anode Zn ZnSO4 CuSO4 Cu Cathode (+) Flow of electric current reactions: At Anode: Zn Zn e - At cathode: Cu e - Overall Reaction: Zn + Cu +2 Cu Zn +2 + Cu representation: Zn(s) ZnSO4 (aq) (1M) CuSO4 (aq) (1M) Cu(s) 8

9 Electro Motive Force The flow of electric current in Galvanic cell is due to potential difference between two electrodes which is known as electrode potential. Electrode potential is a measure of tendency of an electrode to gain or lose electrons. Each electrode has a different tendency to lose or gain electrons. Reduction potential (conventionally used) is the measure of tendency of an electrode to lose electrons and oxidation potential is the measure of tendency of an electrode to lose electrons. Electrodes with high reduction potential readily gains electrons and gets reduced. Electrodes with low reduction potential (oxidation potential) readily loose electrons and gets oxidized. The difference between the electrode potentials of the two electrodes constituting a galvanic cell is known as EMF or cell potential (expressed in volts). EMF = Ecell = E o cathode E o anode = E o right - E o left The reduction potential of an electrode measured at 25 o C, 1atm and 1M concentration of electrolyte is called standard reduction potential. The positive value of Ecell indicates that cell reaction is feasible. Concentration cells Definition: The electrical energy in a concentration cell is due to 'transfer of substance' from the solution of high concentration to the solution of low concentration. The concentration cell is made up of two half cells having identical electrodes and identical electrolytes. However the concentrations of the reactive ions at the two electrodes are different. representation: M M n+ (aq) (C1) M n+ (aq) (C2) M Where C1 and C2 are the concentrations of the active metal ions in contact with the two electrodes respectively. Eg. Silver concentration cell When a metal electrode is dipped in solution containing M n+ ions, then a potential E is developed at the electrode. The value of E varies with the concentration (C) of the ions in accordance with Nernst equation RT nf o E = E + logc 9

10 Electrolytic concentration cell These cells consist of two identical electrodes immersed in two solutions of the same electrolyte at different concentrations. The electrodes of such cells are reversible to one of the ions of the electrolyte. Eg. Zinc concentration cell: Zn Zn +2 (aq) (C1) Zn +2 (aq) (C2) Zn where C1 and C2 are the concentration of Zn +2 at anode and cathode respectively (C2 C1). Thus EMF so developed is due to the transference of metal ions from the solution of higher concentration (C2) to the solution of lower concentration (C1) E = E - E Cathode Anode o 2.303RT o 2.303RT E = E n+ + logc 2 - [E n+ + logc M /M M /M 1] nf nf 2.303RT 2.303RT E = logc 2 - logc1 nf nf 2.303RT E = [logc 2 - logc 1] nf 2.303RT C 2 E = log nf C C 2 E = log (At 298K) n C1 If C1 = C2, then E cell = 0. Applications of concentration cells: 1. To determine the solubility of a sparingly soluble salts. 2. To calculate the valency of cations. 3. To calculate the extent of corrosion in metals. 4. To determine the transition point. 10

11 Batteries The term battery is generally used for two or more galvanic cells connected in series. Thus a battery is a series of portable electrochemical cells which are capable of generating electrical energy. Batteries are of three types: (1) Primary cell (2) Secondary cell (3) Fuel cell. Applications of batteries: The portability of electronic equipments in the form of handsets has been made possible by batteries. A variety of electronic gadgets with more reliability and service have been made more useful with the introduction of rechargeable storage batteries. Comparison between primary, secondary and fuel cells Sl. Primary cells Secondary cells Fuel cells No. 1. reaction is irreversible reaction is reversible. Energy can be withdrawn indefinitely as long as outside supply of fuel is maintained. 2. Must be discarded after use. May be recharged. Cannot be reused. 3. Have relatively short shelf life. Have long shelf life. Have life as long as the fuel is supplied. 4. Function only as galvanic cells. Functions both galvanic & as electrolytic cell. Functions as only galvanic cell. 5. They cannot be used as storage devices They can be used as energy storage devices (e.g. solar/ thermal energy converted to electrical energy) Do not store energy. 6. They cannot be recharged They can be recharged. They cannot be recharged. 7. Eg. Dry cell, Alkaline cell and Li- battery. Eg. Lead acid storage cell, Ni-Cd battery. Eg: H2-O2 fuel cell and CH3OH-O2 fuel cell Primary Batteries In primary cells the cell reaction is not reversible. No electricity is produced after complete conversion of the reactants to products and the cell becomes dead. These batteries are used as source of DC power. Eg: 1. Dry cell (Leclanche cell), 2. Alkaline cell 3. Lithium cell 11

12 1. Dry (Leclanchi ) - Primary Battery Anode: Zn-container Cathode: Graphite rod Electrolyte: Mixture of MnO2(s), NH4Cl (aq) and ZnCl2(s) to which starch is added to make a paste. Voltage: 1.5 V Construction: The anode of the cell is a Zn-container containing an electrolytic mixture of NH4Cl, ZnCl2 and MnO2 to which starch is added to make a thick paste. The graphite rod which is immersed in the centre of the cell acts as cathode. This is a primary cell without fluid component. The anode of the cell is a Zn-container containing an electrolytic mixture of NH4Cl, ZnCl2 and MnO2 to which starch is added to make a thick paste. The graphite rod which is immersed in the centre of the cell acts as cathode. reactions: Anode: Zn(s) Zn 2+ (aq)+ 2e - Cathode: MnO2(s) + NH4 + (aq) + 2e - MnO(OH) - (s) + NH3 Net Reaction: Zn(s) + NH4 + (aq) + MnO2(s) Zn 2+ (aq) + MnO(OH) - (s) + NH3(g) + Elect.Energy Reaction between NH4 + (from NH4Cl) and OH - formed at cathode evolves NH3 gas which disrupts the current flow. This is prevented by the reaction of NH3 (g) with ZnCl2 to form complex [Zn(NH3)2Cl2]. ZnCl2+ 2NH3 [Zn(NH3)2] Cl2 The reactions involved in the dry cell cannot be reversed by passing electricity back through the cell. Hence dry cell is a primary cell. Advantages: Low price; gives voltage of about 1.5 V; normally works without leaking (leak proof cells); possess high energy density; non- toxic; contains no liquid electrolytes. Applications: Dry cell is used in consumer electronic devices like calculators, transistor radios, flash lights, quartz wall clocks, walkman etc. and in small portable appliances where small amount of current is needed. 12

13 2. Alkaline Anode: Zn-container, Cathode: Graphite rod Electrolyte: Mixture of KOH (aq), Zn-powder and MnO2(s) to which starch is added to make a gel. Voltage: 1.5 V Construction: It is an improved form of the dry cell, in which the electrolyte NH4Cl is replaced by KOH. In alkaline battery, the powdered zinc is mixed with KOH and MnO2 to get a gel. A carbon rod (graphite), acts as cathode, is immersed in the electrolyte in the centre of the cell. The outside cylindrical zinc body is made of zinc, act as anode. The EMF of the cell is 1.5 V. reactions: Anode: Zn(s) + 2OH - (aq) Zn 2+ (aq) + 2e - Cathode: 2MnO2(s) + H2O (l) + 2e - Mn2O3(s) + 2OH - (aq) Net Reaction: Zn(s) + 2MnO2(s) + H2O (l) Zn 2+ (aq) + Mn2O3(s) + 2OH - (aq) + Elect.Energy Advantages: 1. Zinc does not dissolve readily in a basic medium. 2. The life of alkaline battery is longer than the dry cell, because there is no corrosion on zinc. 3. Alkaline battery maintains its voltage, as the current is drawn from it. Applications: Alkaline cell is used in consumer electronic devices like calculators, transistor radios, flash lights, quartz wall clocks, walkman etc. and in small portable appliances where small amount of current is needed. 3. lithium - Primary Battery Anode: Lithium metal, Cathode: Metal oxide or TiS2, Electrolyte: Solid polymer and Voltage: 3.0 V Construction: Lithium cell is a primary cell, Lithium metal acts as anode. Metal oxide or TiS2 acts as cathode. Solid polymer will act as an electrolyte; this polymer is packed in between the electrodes. The electrolyte permits the passage of ions but not that of electrons. The EMF of the cell is 3.0V reactions: Anode: Li(s) Li + + e - Cathode: TiS2(s) + e - TiS2 - Net Reaction: Li(s) + TiS2(s) Li + + TiS2 - + Elect.Energy 13

14 Advantages: 1. In this cell voltage is high, 3.0V. 2. Lithium is a light weight metal (7gms). 3. Lithium has the most negative E value. 3. In Lithium cell all constituents are in solid form so that there is no leakage. 4. The battery can be made in a variety of sizes and shapes. Applications: 1. Button sized Lithium batteries are used in calculators, watches, cameras, mobile phones, laptop computers, etc., Secondary Batteries The cells in which the cell reaction is reversed by passing direct current in opposite direction i.e. can operate both as a voltaic cell and as an electrolytic cell. The secondary batteries can be used through a large number of cycles of discharging and charging. They are used as a source of DC power. Eg. Lead acid storage cell, Ni-Cd battery and Lithium ion batteries. 1. Lead acid storage cell - Secondary Battery it Anode: Pb(s), Cathode: PbO2(s), Electrolyte: 20% of H2SO4 (aq) with density of 1.20g/cm 3 Voltage: 6.0 V for 3 pairs, 12.0 V for 6 pairs Construction: It consists of lead antimony alloy coated with lead dioxide (PbO2) as cathode and spongy lead as anode. The electrolyte is a 20% solution of H2SO4. The storage cell can operate both as voltaic cell and electrolytic cell. It acts as voltaic cell when supplying energy and as a result eventually becomes rundown. The cell operates as electrolytic cell when being recharged. The cell consists of a series of Pb-plates (negative plates) and PbO2 plates (positive plates) connected in parallel. The plates are separated from adjacent one by insulators like wood, rubber or glass fiber. 14

15 reactions Discharging (voltaic cell): During the discharging the battery, H2SO4 is consumed the density of H2SO4 falls to 1.20g/cm 3, then battery needs recharging. In discharging, the cell acts as a voltaic cell where oxidation of lead occurs. In charging, the cell is operated like an electrolyte cell and electric energy is supplied to it. At Anodic (-): Pb(S) Pb e - Pb 2+ + SO4 2- (aq) PbSO4(S) At Cathodic (+): PbO2(S) + 4H + + 2e - Pb 2+ +2H2O Pb 2+ + SO4 2- (aq) PbSO4(S) Net reaction: Pb(S) + PbO2(S) + 4H + + 2SO4 2- (aq) 2PbSO4(S) +2H2O + Elect.Energy Charging (electrolytic cell): During this process, lead is deposited at the cathode, PbO2, is formed at the anode and H2SO4 is regenerated in the cell. At Anodic (+): PbSO4(S) + 2H2O PbO2(S) + 4H + + SO4 2- (aq)+ 2e - At Cathodic (-): PbSO4(S) + 2e - Pb(S) + SO4 2- (aq) Net reaction: Elect.Energy + 2PbSO4(S) + 2H2O Pb(S) + PbO2(S) + 4H + +2SO4 2 (aq) Applications: The lead storage cells are used to supply current for electrical vehicles, gas engine ignition, telephone exchanges, electric trains, mines, laboratories, hospitals, broadcasting stations, automobiles and power station. 2. Nickel-Cadmium cell - Secondary Battery Anode: Cd (s), Cathode: NiO2 (s), Electrolyte: 20-28% aqueous KOH jelled with a jelling agent Voltage: 1.4 V Construction: This battery consists of Cd anode and cathode is composed of a paste of nickel dioxide [NiO2]. These cells are rechargeable and give a voltage of 1.4V each. They can be connected in series to give Ni-Cd battery. Electrolyte is 20-28% aqueous KOH jelled with a jelling agent. Discharging (voltaic cell): Anode(-): Cd(s) + 2OH - (aq) Cd(OH)2(s) + 2e - Cathode(+): NiO2(s) + 2H2O + 2e - Net Reaction: Cd(s) + NiO2(s) + 2H2O Ni(OH)2(s) + 2OH - (aq) Cd(OH)2(s)+ Ni(OH)2(s)+ Elect.Energy 15

16 Charging (electrolytic cell): Cathode(-): Cd(OH)2(s) + 2e - Cd(s) + 2OH - (aq) Anode(+): 2Ni(OH)2(s) + 2OH - (aq) NiO2(s) + 2H2O +2e - Net reaction: Elect.Energy + Cd(OH)2(s) + 2Ni(OH)2(s) Cd(s) + NiO2(s) + 2H2O Applications: 1. It is used - in flash lights, photoflash units and portable electronic equipments. 2. In emergency lighting systems, alarm systems. 3. In aircrafts and space satellite power systems. 4. For starting large diesel engines and gas turbines etc. 3. Lithium ion cell - Secondary Battery Anode: Layers of porous carbon, Cathode: layers of lithium metal oxide, Electrolyte: Polymer gel Voltage: 3.8 V 4.2V Construction: It is rechargeable secondary cell. In lithium battery consists of a layers of porous carbon acts as anode and layers of lithium metal oxide acts as cathode. Electrolytes, generally a polymer gel, act as separator between the electrodes. The separator permits the passage of ions but not that of electrons. reactions during discharging (voltaic cell): During discharging, the Li + ions flow back through the electrolyte form negative electrode to the positive electrode. Electrons flow from the negative electrode to the positive electrode. The Li + ions and electrons combine at the positive electrode and deposit there as Li. Li1-xCoO2 + CLix + E.E LiCoO2 + C + Elect.Energy reactions during charging (electrolytic cell): During charging, the Li + ions flow from the positive electrode to negative electrode through electrolyte. LiCoO2 + C + Elect.Energy Li1-xCoO2 + CLix 16

17 Advantages: 1. Lithium ion batteries are high voltage and light weight batteries. 2. It is smaller in size. 3. It produces three times the voltage of Ni-Cd batteries. Applications: 1. It is used in cell phone, note PC, portable LCD TV, semiconductor driven audios. Fuel cells A fuel cell is a galvanic cell in which chemical energy of a fuel oxidant system is converted directly into electrical energy in a continuous electrochemical process. The chemical energy provided by the fuel cell is stored outside the cell. The fuel and oxidants are continuously and separately supplied to the electrodes of the cell where they undergo reactions. Fuel cells are capable of supplying current as long as reactants are supplied. Eg. Hydrogen - oxygen fuel cell and Methanol - oxygen fuel cell 1. Hydrogen Oxygen Fuel cell Anode & Cathode: Porous graphite electrodes impregnated with finely divided Pt/Pd Electrolyte: 25% KOH held in asbestos matrix Voltage: 0.8 V 1.0V Construction: Both anode and cathode consists porous graphite electrodes impregnated with finely divided Pt/Pd. The electrolyte is 25% KOH held in asbestos matrix. The reactants (fuel + oxidant) are constantly supplied from outside and the products are removed at the same rate as they are formed. The product discharged is water and the standard EMF of the cell is 0.8 V 1.0V. reactions: Anode: 2H2(g) + 40H - (aq) 4H2O(l) + 4e - Cathode: O2(g) + 2H2O(l) + 4e - 4OH - (aq) Net Reaction: 2H2(g) + O2(g) 2H2O(l) + Elect.Energy Applications: 1. It is used as energy source in space shuttles e.g. Apollo spacecraft. 2. Used in small- scale applications in submarines and other military vehicles. 3. Suitable in places where environmental pollution and noise are objectionable. 4. Drinking water source in space crafts, submarines and other military applications. 17

18 2. Methyl Alcohol Oxygen Fuel cell Anode & Cathode: Nickel sheet covered by Pt-Pd catalyst Electrolyte: 25 % of KOH solution Voltage: 1.23V Construction: The cell consists of methanol and used as the fuel and oxygen (or) air as the oxidant. Anode: Nickel sheet covered by Pt-Pd catalyst, Cathode: Silver on nickel sheet Electrolyte: KOH solution. In this cell methanol and oxygen are fed to the electrolyte. CO2 produced is absorbed by the electrolyte KOH and converted in carbonates. Formation of carbonates decreases the cell efficiency because of the increasing, concentration and polarization at the electrode surface and decreases the conductivity of the electrolyte. reactions: At Anode: CH3OH(l) + 6OH CO2(g) + 5H2O (l) + 6e At Cathode: 3/2O2(g) +3H2O (l)+ 6e 6OH (aq) Net Reaction: CH3OH(l) +3/2O2 CO2(g) +2H2O (l) + Elect.Energy Advantages of methyl alcohol-oxygen fuel cell: 1. Easy to transport. 2. Do not require complex steam reforming operations. 3. These fuel cells are targeted to portable applications 4. It is excellent fuel due to presence of high concentration of hydrogen in methanol. 5. There is zero emission by the cells hence the fuel cells are eco friendly. Applications: 1. Because of light weight these fuel cells are preferred for space crafts and product H2O is a valuable fresh water source for astronauts. 2. The major application of methyl alcohol-oxygen fuel cells is a fuel for fuel cell motor vehicles like NECAR-5 in Japan, USA. Advantages of Fuel cells over other cells 1. High efficiency in energy conversion. 2. The product H2O is a drinking water source for astronauts. 3. Noise and thermal pollution are low. 4. Low maintenance cost. 5. It is an energy storage system for space applications. 18

19 Important questions 1.Define a reference electrode and explain Construction and functioning of standard hydrogen electrode. 2. Explain Construction and functioning of Calomel electrode. 3. Explain with a neat diagram the function of Reference and ion Selective electrodes? 4. What is electrochemical cells and explain construction and functioning of Galvanic cell. 5. Explain Nernst equation and write its applications. 6. What is electrochemical series and write its applications. 7. Define concentration cells? Explain Electrolyte concentration cell in detail. 8. A zinc road is placed in 0.1M-ZnSO4 solutions at 298 K. write the electrode reaction and calculate the potential of the electrode. Zn = V? 9. What are primary batteries and explain construction and functioning of Dry/Leclanchi battery. 10. Explain Construction and functioning of alkaline cell. 11. Why primary lithium battery best with other primary batteries and explain construction and functioning of lithium battery. 12. What are secondary batteries and explain construction and functioning of Lead acid-storage battery. 13. Explain Construction and functioning of Ni-Cd cell. 14. Why secondary lithium ion batteries are best with other secondary batteries and explain construction and functioning of lithium ion battery. 15. What is a fuel cell and explain Hydrogen Oxygen Fuel cell. 16. What are the advantages of fuel cells over primary and secondary batteries and Explain Methanol Oxygen Fuel cell. Objective questions 1. Which of the following is weak electrolyte [ ] (a) HCl (b) NaOH (c) NH4OH (d) NaCl 2. Pure water does not conduct electricity because, it is [ ] (a) Acidic (b) Basic (c) Almost unionized (d) Decomposed easily 3. The potential of standard hydrogen electrode is taken as [ ] (a) 1 volt (b) 0 volt (c) 10 volts (d) None of these 19

20 4. Which statement is correct about the Daniel cell [ ] (a) The spontaneous cell reaction involves the reduction of Cu +2 to Cu. (b) The spontaneous cell reaction involves the oxidation of Cu by Zn. (c) The spontaneous cell reaction involves the reduction of Zn +2 to Zn. (d) The Daniel cell is an example of an electrolytic cell. 5. A galvanic cell converts [ ] (a) Electrical energy into chemical energy. (b) Chemical energy into electrical energy. (c) Electrical energy into heat energy. (d) Chemical energy into heat energy 6. Electrochemical series the elements are arranged in the [ ] (a) Decreasing order of standard reduction potentials. (b) Increasing order of standard reduction potentials. (c) Increasing order of equivalent weight. (d) Increasing order of oxidation potentials. 7. The ion selective electrodes make use of [ ] (a) Membrane (b) Mercury (c) KCl solution (d) Platinum electrode 8. The electrode potential is the tendency of a metal [ ] (a) To gain electrons. (b) To loss electrons. (c) Either to gain or loss electrons. (d) None of the above 9. Calomel electrode is constructed using a solution of [ ] (a) Saturated KCl (b) Saturated CaCl2 (c) Saturated MgCl2 (d) Saturated NaCl 10. A storage cell is a device that can operate [ ] (a) Both as voltaic cell and electrolytic cell. (b) As voltaic cell (c) As electrolytic cell. (d) As concentration cell 11. Which one of the following is a primary cell? [ ] (a) Ni-Cd cell (b) Lead-Acid cell (c) lithium ion cell (d) Alkaline cell 20

21 12. The cathode of Ni-Cd battery is composed of [ ] (a) Cd (b) Ni (c) Paste of NiO(OH) (d) Paste of Cd (OH)2 13. A fuel cell converts [ ] (a) Chemical energy of fuels directly to electrical energy. (b) Chemical energy of fuels directly to heat. (c) Chemical energy of fuels directly to pressure. (d) Electrical energy of fuels directly to chemical energy. 14. What is the voltage produced by the H2 - O2 fuel cell. [ ] (a) 1.0 V (b) 1.23 V (c) 2.0 V (d) 0.5 V Fill in the blanks 1. The substance which allows electric current to pass through it is called. 2. The substance which conducts electricity without decomposition is called. 3. In an electro chemical cell electron flows from to. 4. are the cells which do not store the energy. 5. Electro chemical series is based on. 6. A device which converts electrical energy to chemical energy is called. 7. The substance which conducts electricity without decomposition is called. 8. The standard electrode potential of saturated calomel electrode at 25 0 c is. 9. When two identical electrodes are dipped in two electrolytic solutions of different concentrations, the resulting obtained is called. 10. A cell whose reaction is not reversible is called. 11. In lithium primary cell, the anode is. 12. In alkaline cell, electrolyte is. 13. Cathode in dry cell is. 14. The electrolyte in a Lead-acid battery is. 15. In lead-acid storage battery during discharging operation the concentration of H2SO In Ni-Cd cell the cathode is. 21

22 Two Marks and Three Marks Questions 1. Define single electrode potential mention the factors affecting it. 2. Define reference electrode give one example. 3. Write the construction of standard Hydrogen electrode 4. Write the construction of saturated calomel electrode 22

23 5. Write the Nernst equation. 6. What is electro chemical series? Write its applications. 7. What is electro chemical cell give one example? 8. What is Concentration cell. 23

24 9. What is EMF and calculate the potential of a zinc road is placed in 0.1M-ZnSO4 solutions at 298 K. write the electrode reaction and calculate the potential of the electrode. Zn = V? 10. Write the chemical reactions involved in Methanol-Oxygen Fuel cell. 11. Write the differences between Primary and Secondary batteries 12. Write Discharge and charging reactions of lead-acid battery. 24

25 13. Write Discharge and charging reactions of Ni-Cd battery. 1. Standard Hydrogen Electrode. Draw the neat diagrams of the following 2. Calomel electrode: 25

26 3. Glass electrode: 4. Daniel cell: 5. Electrolytic Concentration cell: 26

27 6. Dry / Leclanche : 7. Alkaline : 8. Lithium Battery: 27

28 9. Lead-Acid Storage Battery: 10. H2-O2 fuel : 11. CH3OH-O2 fuel :: 28