Electrochemistry: Conductors, Cells & Batteries PDF

ELECTROCHEMISTRY Electrochemistry is a branch of chemistry, which deals with the relationship between electrical energy and chemical changes taking place in redox reactions. i.e., how chemical energy or how electrical energy can be used to bring about a redox reaction which is otherwise not spontane...

ELECTROCHEMISTRY Electrochemistry is a branch of chemistry, which deals with the relationship between electrical energy and chemical changes taking place in redox reactions. i.e., how chemical energy or how electrical energy can be used to bring about a redox reaction which is otherwise not spontaneous. It has many applications in electrolysis, energy producing cell etc. ❖ Conductors and non-conductors: The substances which conduct electricity are called conductors. For eg: Metals – Cu, Ag, Fe, Au Salts – NaCl, CdS, PbS etc Acids – HCl, H2SO4, HNO3 Alkali – NaOH, KOH, Ca(OH)2 The substances which cannot conduct electricity are called non-conductors. For eg: sugar, glucose, P, S etc. ❖ Classification of conductors: a. Metallic conductors: The conductors through which conduction of electricity takes place by the migration of electrons under the influence of an applied potential is called metallic conductors. For eg: metals, alloys, graphite etc. b. Electrolytic conductors: The conductors through which the conduction of electricity takes place by the migration of ions towards oppositely charged electrodes due to occurrence of chemical changes at the surface of electrodes are called electrolytic conductors. For eg: solution of weak and strong electrolyte, molten salt etc. ❖ Strong and weak electrolyte: Strong electrolyte Weak electrolyte The electrolytes which ionize or The electrolyte which ionize to dissociate almost completely small extent in aqueous into ions in aqueous solution solution are called weak are called strong electrolyte. electrolytes. The solution of strong Aqueous solution of weak electrolyte is a good conductor electrolyte is a poor conductor of electricity and has a high of electricity and has low value value of covalent conductance of equivalent conductance. Eg: even at low concentrations. Eg: Acids: H2SO3, H2CO3. Acids: HCl, H2SO4, HNO3. Bases: Bases: NH4OH. NaOH, KOH. ❖ Arrhenius Theory of ionization(Svante Arrhenius-1884): The basic postulates of Arrhenius Theory of ionization are as follows; 1. When an electrolyte is dissolved in water then it splits into two types of charged particles called ions. The positively charged particle is called cation and negatively charged particle is called anion. This process is called ionization. NaCl ↔ Na+ +Cl- 2. The electrolytic solution is electrically neutral. i, e. the total number of positively charged cations and negatively charged anions are equal. H2SO4(aq) ↔ 2H+(aq) + SO4-2(aq) 3. The process of ionization is reversible. The ions present in the solution are constantly reuniting and molecules are dissociating. Thus, there is a state of dynamic equilibrium between the undissociated molecules and ions. For eg: H2SO4(aq) ↔2H+(aq) + SO4-2(aq) NaCl ↔ Na+ +Cl- 4. The electrical conductivity of an electrolyte is due to the migration of ions towards oppositely charged electrodes. i, e. cations move towards cathode and anions move towards anode. 5. The degree of ionization gives the extent of ionization of an electrolyte. i, e. Total no.of ions produced Degree of ionization(α) = Total no.of unionized molecules 6. The dissociation of electrolyte depend upon; -Nature of electrolyte -Degree of dilution - Temperature 7. The electrical conductivity depends upon the number of ions present in the solution and speed of ions. ❖ Electrochemical cell: An electrochemical cell is an arrangement of two electrodes either in same electrolytic solution or in different electrolytic solution which is capable of producing electricity due to chemical reaction within the cell or chemical reaction due to passage of electricity. On the basis of function, electrochemical cells are of two types; 1. Electrolytic cell 2. Galvanic (Voltaic) cell 1. Electrolytic cell: The electrochemical cell in which the electrical energy is used to bring chemical reaction is called electrolytic cell. In this cell, two metallic electrodes are dipped into the solution of suitable electrolyte. Then, the electrodes are connected to the external source of electricity such as battery. The electrode which is connected to the positive terminal of the battery is called anode and the electrode which is connected to the negative terminal of the battery is called cathode. When electricity is passed, chemical reactions takes place inside the cell. i,e. cations move towards  cathode and anions move towards anode. For eg; Let us consider the electrolysis of molten NaCl. Reactions involved: NaCl ↔ Na+ + Cl- (Molten) At anode: Cl- - 1e- → Cl (oxidation) Cl + Cl → Cl2 At cathode: Na+ + 1e- → Na ( reduction) 2. Galvanic/Voltaic cell: The electrochemical cell in which electrical energy is produced from chemical reactions are called galvanic or voltaic cell. In voltaic cell, oxidation and reduction reaction occur in separate compartment called half-cell. The spontaneous reaction is responsible for the production of electrical energy in galvanic cell/voltaic cell. For eg: Daniel cell Daniel cell: It is a simple zinc – copper cell which is formed by the combination of zinc electrode dipped in the solution of its salt ( eg; 1M ZnSO4 ) and copper electrode dipped in the solution of its salt (eg; 1M CuSO4). These two electrolytes are taken in two different beakers and two electrodes are joined externally by means of a wire to the voltmeter and internally by means of a salt bridge. A salt bridge is U-shaped glass tube containing the saturated solution of an electrolyte like KCl, KNO3 or NH4NO3 in agar-agar or gelatin. Agar-agar is a natural vegetable gelatin counterpart. It is white and semi-translucent when sold in packages as washed and dried strips or in powdered form. It can be used to make jellies, puddings, and custards. When making jelly, it is boiled in water until the solids dissolve. The arrangement of Daniel cell is shown in the figure below; Fig: Daniel cell – an example of galvanic cell In Daniel cell, oxidation half reaction takes place at Zn-electrode called anode and reduction half reaction takes place at Cu-electrode called cathode. Since, the electron moves from zinc electrode to copper electrode, the Zn-electrode is considered as negative terminal (anode) and Cu- electrode is considered as positive terminal (cathode). The flow of electron takes place from negative terminal (anode) to positive terminal (cathode) whereas flow of current takes place from cathode to anode.  Cell reaction: In electrochemical cell, oxidation half and reduction half reaction takes place separately at anode and cathode respectively. The net chemical reaction obtained by adding half reactions is called cell reaction. At anode: Zinc looses two electrons and passes into the solution as Zn++. Zn – 2e_ → Zn++ At cathode: Cu++ gains two electrons and reduced to Cu which deposits at cathode. Cu++ + 2e_ → Cu Overall cell reaction: Zn + Cu++ → Zn++ + Cu Cell notation: Zn(s) | ZnSO4(1M) || CuSO4(1M) | Cu(s) Anode salt bridge cathode ❖ Salt bridge and its functions: Salt bridge is U-shaped glass tube containing a saturated solution of an inert electrolyte having cation and anion with same mobility like KCl, KNO3, NH4NO3 in agar-agar gel or gelatin. Inert electrolyte refers to those electrolytes which do not take part in the redox reaction. The function of salt bridge are; i. It helps to prevent the physical contact between two electrolytic solutions. ii. It helps to complete the electrical circuit by connecting the two solutions of half cells. iii. It helps to maintain electrical neutrality in the compartment. ❖ Rules for writing cell notation: The representative of galvanic/voltaic cell is done as follows; 1. A single vertical line (|) is used to separate metal electrode and ion in electrolytic solution. For eg; Zn|Zn++(aq) Cu++(aq)|Cu(s) Oxidation half-cell Reduction half-cell 2. The anode half-cell is always written in the left but cathode half-cell is written in the right. 3. A double vertical line (||) represents the salt bridge preventing the electrolytic solution from mixing, is written in the middle of two half cells. For eg; Zn(s)|Zn++(aq) || Cu++(aq)|Cu(s) 4. The concentration of the electrolytic solution is placed inside the bracket. For eg; Zn++(1M), Cu++(0.5M) ❖ Single electrode potential and origin of electrode potential: When a piece of metal/electrode (M) is placed into a solution containing one of its salt, then few metal atoms (M) pass into the solution leaving electrons at the surface of the metal/electrode and attains a state of dynamic equilibrium as M ↔ Mn+ + ne- When equilibrium state is reached an electric double layer is formed between oppositely charged species. The electric double layer thus formed is held together by means of electric force of attraction between the metal/electrode and the solution in equilibrium. This force of attraction is called single electrode potential/electrode potential of the electrode system(M/Mn+). The magnitude of electrode potential depends on the intensity of electric double layer which further depends on the nature of electrode and concentration of its solution. Higher the intensity of electric double layer higher will be the electrode potential. It is expressed in hydrogen scale.  Depending upon the capacity of metal(electrode) to loose or gain electrons, electrode potential is of two types; 1. Oxidation potential: The electrode potential produced due to the loss of electrons or oxidation of metal electrode is called oxidation potential. For eg; Zn ↔ Zn++ + 2e- Cu ↔ Cu++ + 2e- 2. Reduction potential: The electrode potential produced due to the gain of electrons or reduction of metal electrode is called reduction potential. For eg: Cu++ + 2e- ↔ Cu Ag+ + 1e- ↔ Ag Oxidation and reduction potential of an electrode are numerically equal but opposite in sign. For example: If the oxidation potential of Cu|Cu++ is -0.34V then the reduction potential of Cu++|Cu will be +0.34V. ❖ Standard electrode potential: The electrode potential measured at standard conditions i,e. 298K temperature, 1 atm pressure and 1 molar concentration of electrolytic solution is called standard electrode potential. The single electrode potential and standard electrode potential of all the electrodes are expressed with respect to the electrode potential of standard  hydrogen electrode whose potential is arbitrarily taken to be equal to 0 volt at all temperature. The oxidation potential measured at 1M concentration, 1 atm. Pressure and 298K temperature is called standard oxidation potential. Whereas, the reduction potential measured at 1M concentration, 1 atm. Pressure and 298K temperature is called standard reduction potential. The electrode having tendency to lose electron greater than standard hydrogen electrode( taken 0 V) will have positive value of oxidation potential and the electrode having tendency to gain electron greater than that of standard hydrogen electrode will have positive value of reduction potential. However, the electrode potential is usually expressed in standard reduction potential. ❖ Factors affecting magnitude of electrode potential: a. Nature of metal/electrode: Highly reactive metals like K, Na, Ca etc. have strong tendency to lose electrons, hence have higher electrode potential. But less reactive metals like Cu, Ag, Au etc. have less tendency to lose electrons, hence have lower electrode potential. b. Concentration of metal ions in the solution: As the concentration of metal ions increases in the solution, electrode potential decreases and vice-versa. c. Temperature: Electrode potential and temperature are directly proportional to each other. i,e. on increasing the temperature of the electrolytic solution, electrode potential also increases. ❖ Electromotive force(emf)/cell potential: The difference between the electrode potential of two half cells (oxidation half and reduction half) is known as emf of cell or cell potential. It is calculated as Ecell = reduction potential - reduction potential of Of cathode anode  Ecell = Ecathode - Eanode …………………………….(i) For standard electrode potential, Eocell = Eocathode - Eoanode …………………………(ii) During the calculation of emf of a galvanic cell, the reduction potential is used. If oxidation potential is given, its sign should be changed to get the reduction potential. ❖ Types of electrodes: The metallic rods dipped into the electrolytic solution in an electrochemical cell are called electrodes. There are two types of electrodes on the basis of electrode potential; a. Indicator electrode: The electrode whose potential is to be determined is called indicator electrode. b. Reference electrode: The electrode whose potential is known is called reference electrode. Reference electrode is of two types; i. Primary reference electrode (Standard hydrogen electrode) ii. Secondary reference electrode (Calomel electrode) i. Standard hydrogen electrode: Standard hydrogen electrode is prepared by passing dry and pure hydrogen gas at constant 1 atm. Pressure over a Platinized Pt gauge attached with a Pt wire which is dipped in an aq. Solution of 1M concentration of H+ ion. The platinum electrode allows electrical contact to be made. Pt absorbs hydrogen on the surface and it ensures good contact between H+ and absorbed hydrogen so that electrode reaction occur. But being an inert metal, Pt doesn’t take part in reaction. The platinum electrode needs to be coated with finely divided Pt called platinized platinum. The reactions involved are; a. H + e- → 1/2H2 ; If another half cell gives e- to SHE. + b. 1/2H2 → H+ + e- ; If standard hydrogen electrode gives e- to other half cell.  Therefore, SHE can acts as cathode or anode relative to another half cell. The notation of SHE is; Pt/H2(g) | H+(aq) ( 1 atm) (1M) Under these conditions(1 atm pressure and 1M concentration), the electrode potential of SHE is found to be 0.00V at all temperature. Fig: Standard hydrogen electrode (SHE) Uses of SHE: The standard hydrogen electrode(SHE) is used to to measure the electrode potential of all electrodes by combining the given electrode system with standard hydrogen electrode. Emf of the cell is calculated as; Ecell = EMn+/M - EH+/H2 = EMn+/M - 0 = EMn+/M Disadvantages of SHE: Standard hydrogen electrode has some limitations in its construction and use; 1. It is difficult to obtain pure H2 gas. 2. It is difficult to maintain 1 atm. Pressure of H2 gas throughout the experimental measurement. 3. It is difficult to maintain ideally pure Pt electrode because it gets corroded or poisoned. When it is poisoned it is to be platinized repeatedly for fresh use. 4. Electrode is not easy to transport. ii. Calomel electrode: Standard hydrogen being not handy, it is replaced by a secondary standard electrode known as calomel electrode. To construct calomel electrode, a glass tube sealed with Pt wire at the bottom is taken and within the tube it is dipped in mercury layer and the layer is covered with a paste of Hg and Hg2Cl2. The remaining upper layer of the tube is filled with either a normal (1N) or saturated KCl solution. Fig: saturated calomel electrode (SCE) Reaction involved are: Hg2Cl2 + 2e- ↔ 2Hg + 2Cl- The notation of SCE is (Pt)/Hg(l) |Hg2Cl2(s)/Cl-(aq)  The electrode reaction is reversible and electrode potential of calomel electrode depends on the concentration of Cl- and temperature. For saturated calomel electrode(SCE), the reduction potential at 25oC is found to be 0.2415V. If a normal calomel electrode having 1M Cl- solution is used, the reduction potential at 25oC is found to be 0.2800V. Advantages of calomel electrode: a. It is easy to set up and attains equilibrium of electrode process rapidly. b. It is compact and of smaller size so that it can be easily transported. c. It can be connected to salt bride easily. d. It is a reversible electrode and it work as anode or cathode. e. It provides constant potential throughout the experiment. Practice: Q. What are the basic differences between electrolytic cell and voltaic cell? Q. What is the function of platinized Pt in SHE? Write the half reaction when SHE acts as cathode and anode. Q. Define indicator and reference electrode with example. Q. Define: a) Standard electrode potential b) Oxidation potential c) Reduction potential ❖ Electrochemical series or activity series: If some electrode systems are arranged in increasing order of their standard reduction potential value then a vertical series of the elements is obtained which is known as electrochemical series or activity series. The electrochemical series with some electrode system and their standard reduction potential is shown below; ❖ Applications of electrochemical series: 1. To compare oxidizing and reducing power: The substance having lower standard reduction potential value is a strong reducing agent. Whereas, the substance having higher standard reduction potential value is a strong oxidizing agent. For eg: Zn++ + 2e- → Zn Eocell = -0.76V Cu++ + 2e- → Cu Eocell = +0.34V Hence, Cu++ is the strong oxidizing agent than Zn++. 2. To construct galvanic cell: The electrode system having higher value of reduction potential acts as a cathode and the electrode system having lower reduction potential acts as an anode of the galvanic cell to be constructed. For eg: In Al-Ag cell, If standard reduction potential of Al-electrode is -1.66V and that of Ag-electrode is +0.80V. Hence, Al-electrode acts as anode but Ag-electrode acts as cathode. 3. To calculate emf of the cell: EMF of the galvanic cell is calculated by the algebraic sum of oxidation potential of anode and reduction potential of cathode. Ecell = Ered of cathode + Eoxd. of anode Ecell = Ecathode – Eanode………………………(i) 4. To predict spontaneity/feasibility of redox reaction: For a given redox reaction to be feasible, the oxidation potential of reducing agent should be greater than that of oxidizing agent. Alternatively, if the calculated emf of the cell is positive, the reaction will be feasible and spontaneous but if the calculated emf is negative, the reaction will be non-spontaneous and it is not feasible.  5. To predict the relative reactivity of metals: The metal having higher standard reduction potential has low reactivity and vice-versa. For eg; Eo of Li+/Li = -3.05V Eo of Cu++/Cu = +0.34V Hence, copper is less reactive than Li. 6. Idea of displacement reaction: The metal having negative value of reduction potential( less than hudrogen) can displace hydrogen from acids. In other hand, the metals having lower reduction potential (more electropositive) can displace the metals having higher reduction potential (less electropositive) from their salt solution. For eg: Zn + dil. H2SO4 → ZnSO4 + H2↑ Fe + CuSO4 → FeSO4 + Cu↓ Practice: Q. Calculate the standard EMF of a cell which involves the following reaction Zn + 2Ag+ → Zn++ + 2Ag Given that: EoZn++/Zn = -0.76V and EoAg+/Ag = 0.80V Q. Given standard reduction potential are EoZn++/Zn = -0.76V and EoSn++/Sn = -0.14V. Using these values a. Construct a galvanic cell indicating anode and cathode. b. Calculate the standard EMF of the cell. c. Write the cell reaction Q. Predict whether the following reaction is feasible or not at 298K. Co(s) + Fe++(aq) → Co++(aq) + Fe(s) Given, EoCo++/Co = -0.28V EoFe++/Fe = -0.44V Q. Can a nickel spatula be used to stir a solution of copper sulphate? Support your answer with a reason. Given, EoNi++/Ni = -0.25V EoCu++/Cu = +0.34V ❖ EXAMPLES OF GALVANIC/VOLTAIC CELL: a. Zn – Cu cell/Daniel cell: Daniel cell is constructed by the combination of Zn- electrode and Cu-electrode. It is given that the standard reduction potential of the electrodes are EoZn++/Zn = -0.76V and EoCu++/Cu = +0.34V. The standard reduction potential of Cu-electrode is higher than Zn-electrode. Therefore, Cu-electrode acts as cathode and Zn-electrode acts as anode. Cu- electrode is prepared by dipping Cu rod in 1M CuSO4 solution in a beaker and Zn-electrode is prepared by dipping Zn rod in 1M ZnSO4 solution in another beaker. The cell is then constructed by connecting the electrodes internally by means of salt bridge and externally by means of a wire to a voltmeter. The cell is represented by the cell notation as; Zn(s) | Zn++(1M) || Cu++(1M) | Cu(s) Anode salt cathode Bridge  Reactions involved are; At Zn-anode (-ve pole): Zn - 2e- → Zn++ (oxidation half) At Cu-cathode (+ve pole): Cu++ + 2e- → Cu (reduction half) Overall reaction: Zn + Cu++ → Zn++ + Cu This spontaneous cell reaction brings flow of electrons from anode to cathode across the external circuit as a result of which an electric force develops in the direction opposite to the motion of electrons in the external circuit which is called electromotive force (emf) or cell potential. Standard emf/cell potential of Zn-Cu/Daniel cell: EoCell = EoCathode - EoAnode = 0.34V - (-0.76V) = 0.34V + 0.76V = 1.1V b. Ag-Cu cell: Silver-copper(Ag-Cu) voltaic/galvanic cell is constructed by the combination of Ag-electrode and Cu-electrode. It is known that the standard reduction potential of Ag+/Ag = +0.80V and the standard reduction potential of Cu++/Cu = +0.34V. Since the standard reduction potential of Ag-electrode is higher than Cu-electrode, therefore Ag-electrode acts as cathode and Cu- electrode acts as anode. Ag-electrode is prepared by dipping Ag rod in 1M AgNO3 solution and Cu-electrode is prepared by dipping Cu rod in 1M CuSO4 solution. The cell is then constructed by connecting the electrodes internally  by means of salt bridge and externally by means of a wire to voltmeter. The cell is then represented by cell notation as; Cu(s) | Cu++(1M) || Ag+(1M) | Ag(s) Anode salt cathode Bridge Reactions involved: At Cu-anode: Cu -2e- → Cu (oxidation half) At cathode: Ag+ + 1e- → Ag ( reduction half) Overall reaction: Cu + 2Ag+ → Cu++ + 2Ag This spontaneous reaction is responsible to produce cell potential or emf of the Ag-Cu cell. Standard emf/cell potential of Ag-Cu cell: EoCell = EoCathode - EoAnode = +0.80V – (+0.34V) = 0.46V ❖ Relationship between cell potential and free energy: A voltaic or galvanic cell is a device that generates electrical energy in the expense of chemical energy produced from spontaneous reaction occurred in the cell. The chemical energy transformed into electrical energy is measured as decrease of Gibb’s free energy. It means, the electrical work done by the cell is equal to the decrease of Gibb’s free energy of the reaction. Electric work done by cell = quantity of electric charge transported × emf of cell or, Electric work done = nFE where, n = No. of moles of e-s flowing cell F= Faraday’s constant (96500C) E= emf of cell The electric work done by the cell is equal to the decrease of Gibb’s free energy of cell reaction. i,e. -∆G = nFE …………………………………(i) For standard state, -∆Go = nFEo ………………………………..(ii) where, ∆Go = standard free energy change Eo = standard emf of the cell Equation (i) and (ii) are the expressions for the relation between cell potential/emf and Gibb’s free energy change. The change in Gibb’s free energy is given by ∆G = ∆H - T∆S …………………………….(ii) ❖ Commercial Batteries: Generally, the arrangement of few numbers of cells for the same type in a series connection is known as battery. Electrochemical cells/batteries are widely used in vehicles, radio, watch, mobile, telephone etc. Some are batteries are rechargeable and some are not rechargeable. Rechargeable battery needs electricity to recharge them. However, every electrochemical cell/ battery isnot suitable for commercial purpose. A commercial battery should have the following characteristics; a. It should be light and must be of compact in size. b. It should give a constant voltage during its use.  The commercial battery is of two types; 1. Primary cells 2. Secondary cells 1. Primary cells(Non-rechargeable cells/Batteries): In primary cells, the reaction occurs in only one direction and cannot be reversed. As a result, primary cells become dead after using over a period of time. Therefore, it cannot be charged or reused due to irreversible nature of reaction. For eg; Dry cell, Fuel cell, Cadmium cell etc a. Dry cell(Leclanche cell/ Zn - C battery): The most common type of primary cell which is a compact form of Leclanche cell developed by a French engineer George Leclanche in 1866 is known as dry cell/Leclanche cell. It is widely used in portable electric and electronic equipments. Each cell has cell voltage of 1.5V. It consists of an outer container made up of zinc which acts as anode and provides electrons to the external circuit by the anode reaction. The zinc is lined from inside with an insulating paper. The carbon rod having a brass cap acts as a cathode. The space between cathode and anode is filled with a moist mixture of MnO2, thick paste of NH4Cl(electrolyte), ZnCl2(as deliquescent) and charcoal. The porous paper lining prevents direct contact between zinc container and the paste. This acts as a salt bridge. The cell is sealed from the top with wax. Reactions involved are; At anode: Zn → Zn++(aq) + 2e- Zn++ then migrates towards carbon electrode (cathode). At cathode: MnO2(s) + H2O(l) + e- → MnO(OH)(s) + OH-(aq) Then OH- migrates towards anode and combines with NH4+ as, NH4+ + OH- → NH3 + H2O Overall reaction at cathode, MnO2(s) + NH4+ + e- → MnO(OH)(s) + NH3 The trace reaction occurs at cathode is; 2NH4+ + 2e- → 2NH3 + H2 Where H2 is immediately trapped by MnO2 and forms Mn2O3 and H2O in traces. In cathode reaction, NH3 is not liberated as a gas but combines with Zn++ coming from anode to form Zn(NH3)42+ ion. Zn++ + 4NH3 → [Zn(NH3)4]2+ The formation of stable complex ion Zn(NH3)42+ lowers the concentration of free Zn++ and increases the voltage of the cell. The dry cell has a potential of about 1.5V. At the end of its life, the concentration of Zn(NH3)42+ increases and its chloride salt crystalizes out. The salt can be partially diffused away from the anode by warming an exhausted cell gently for a few hours. In this way, the potential of cell can be partially restored. Dry cell cannot be recharged. It has no definite life due to the acidic nature of NH4Cl which corrodes Zn container continuously even when not in use. Fig: Zinc – Carbon dry cell  Q. Why is dry cell not rechargeable? b. Fuel cells: Fuel cells are primary cells in which reactants are continuously supplied to the electrodes of the cells from outside as fuels. The electrical cells that convert the energy from the combustion of fuels such as hydrogen, carbon monoxide or methane directly into electrical energy are called fuel cells. Such cells produce voltage of 1.2V. For eg: Hydrogen – Oxygen fuel cell Fig: Hydrogen – Oxygen fuel cell In hydrogen-oxygen fuel cell, hydrogen and oxygen are bubbled through a porous carbon electrode into concentrated aq. NaOH. Catalysts are incorporated in the electrode. Reactions involved in this cell are; At anode: [H2(g) + 2OH-(aq) → 2H2O(l) + 2e-] × 2 (Oxidation half) At cathode: O2(g) + 2H2O(l) + 4e- → 4OH- (Reduction half) Overall reaction: 2H2(g) + O2(g) → 2H2O(l) Fuel cells ( Hydrogen – oxygen cell ) run continuously as long as the reactants are supplied. Since, fuel cells convert the energy of fuel directly into electrical energy, they are potentially more efficient and pollution free than conventional method of generating electricity. However, the efficiency of such fuel cells is only about 60-70 percent. Fuel cells are widely used in transportation, material handling and stationary, portable and emergency backup power, space mission, heart pace maker etc. Advantage: a. They are potentially more efficient. b. They are pollution free. Limitations of hydrogen-oxygen fuel cells: a. Hydrogen is very difficult to store safely in a vehicle. b. Large amount of hydrogen are not readily available. ❖ Secondary cells (Rechargeable cell/Battery): The secondary cells are can be recharged by passing electricity through them after every use. Hence, such cells can be used again and again. In these cells, the electrical energy is stored in the form of chemical energy. Therefore, they are called as storage or accumulators. They act as galvanic cells during discharge and electrolytic cells during recharging. For examples: Lithium-ion battery, Acid lead storage battery, Nickel-Cadmium storage battery etc. a. Acid lead storage battery: Acid lead storage batteries are being used commercially from long ago. They consist of grid electrodes. i, e. one is a grid of Pb-Sb impregnated with spongy Pb acting as an anode and another is a grid of Pb-Sb alloy impregnated with PbO2 acting as a cathode. These grid electrodes are immersed in a solution of 20% -38% H2SO4 which acts as an  electrolyte. the schematic diagram of this battery is given in the figure below; On discharging, At anode: The spongy lead is oxidized to lead ions and lead plate accumulate a negative charge. Pb → Pb++ + 2e- (Oxidation) The lead ions then combines with sulphate ions of H2SO4 and forms insoluble lead(II) sulphate which begins to coat the lead electrode. Pb++ + SO4-- → PbSO4 (Deposition) Thus, the net reaction at anode during discharge is Pb + SO4-- → PbSO4 + 2e- At cathode: The electron produced at the lead electrode travel through external circuit towards cathode containing PbO2. PbO2, in the presence of H+ ions is reduced to Pb++ ions which combine with SO4—of H2SO4 to form PbSO4 that coated on PbO2 electrode. PbO2 + 4H+ + 2e- → Pb++ + 2H2O (Reduction) Pb++ + SO4-- → PbSO4 Thus, the net reaction at cathode during discharge is; PbO2 + 4H+ + SO4-- + 2e- → PbSO4 + 2H2O Hence, overall reaction during discharge is; Pb + PbO2 + 4H+ 2SO4-- → 2PbSO4 + 2H2O ECell = +2.041V  If the density of sulphuric acid falls below 1.20 g/cc then the cell must be recharged. During recharge, the cell reaction takes place in reverse direction due to external power supply. On recharging, the reaction involved are; At anode: PbSO4 + 2H2O → PbO2 + 4H+ SO4-- + 2e- At cathode: PbSO4 + 2e- → Pb + SO4— Hence, overall reaction during recharging is; 2PbSO4 + 2H2O → Pb + PbO2 + 2H2SO4 The battery after full charge can be stored upto one year safely and the life of this type battery is 12+ years. b. Lithium – ion battery(Lithium-polymer): Lithium-ion battery is also an example of secondary cell/rechargeable battery. The commercial use of this battery has been started after 1990 A.D. The cell consists of anode(-ve electrode), cathode(+ve electrode), semi- permeable barrier(separater) and electrolyte solution in non-aqueous solvent. The anode is the graphite which is lithiated/intercalated as LiC6. It produces Li+ ions while discharging and stores lithium while charging. The cathode is either LiCoO2 , LiFePO4 and or LiNiMnCoO2 (NMC). And the electrode present between the electrodes is lithium salt like LiClO4, LiBF4 etc. prepared in non-aqueous organic solvents like ethylene carbonate. Handheld Electronic Company prepared the Lithium-ion battery by using LiCoO2 material as cathode, LiC6 as anode and polymer gel electrolyte. The cell is represented as; Lithiated nLi-polymer gel or LiC6 || LiClO4 solution in || LiCoO2/LiFePO4 Ethylene carbonate Anode Electrolyte cathode In this cell, lithium undergoes oxidation to Li+ ions which move from anode to cathode where Li+ gets stored during discharge and the reactions reverses during charging. The semi-permeable barrier allows Li+ ion to move from  anode to cathode and vice-versa. The barrier prohibits the flow of electrons inside the structure of the battery. Electrons are allowed to move across the external circuit generating cell potential of the magnitude 3.6V to 3.8V. The reactions involved are; At anode: LiC6 - 1e- → C6 + Li+ (Oxidation) At cathode: CoO2 + Li+ + 1e- → LiCoO2 ( Reduction) Overall reaction: LiC6 + CoO2 ↔ C6 + LiCoO2 In general, lithium ion battery has high energy density(250-690Wh/L), low self-discharge rate(0.35-2.5% per month). The charge and discharge efficiency of lithium-ion battery is 80% - 90% and rechargeable duration is 400-1000 cycles. Fig: Lithium-ion battery Uses of Lithium ion battery: a. Li-ion battery are used in portable electronic devices such as laptop, digital camera, mobile, flash light, medical equipments. b. They are also used in electric vehicles. c. They are also used in electric invertors, aerospace, military etc.  Disadvantages: a. They cannot be used in micro-oven, pressure container and other devices in which heating is done as it contains volatile flammable solvent.