Electrochemistry Notes 2024-25 PDF

ELECTROCHEMISTRY Scope : (i) Electrochemical cells: introduction, redox reactions (principle of oxidation and reduction in a cell). (ii) Galvanic cells - introduction; representation, principle – oxidation reduction. Mechanism of production of electric current in a galvanic cell....

ELECTROCHEMISTRY Scope : (i) Electrochemical cells: introduction, redox reactions (principle of oxidation and reduction in a cell). (ii) Galvanic cells - introduction; representation, principle – oxidation reduction. Mechanism of production of electric current in a galvanic cell. o (iii) Measurement of potential. Single electrode potentials.Standard hydrogen electrode (E ) - definition, preparation, application and limitations.Standard electrode potential - Measurement ++ ++ of standard electrode potential of Zn / Zn, Cu / Cu, half cell (using standard hydrogen electrode).Cell notation – representation. Factors affecting electrode potential with explanation - main emphasis on the temperature, concentration and nature of the electrode. (iv) Electrochemical series. Its explanation on the basis of standard reduction potential. Prediction of the feasibility of a reaction. (v) Nernst equation and correlation with the free energy of the reaction with suitable examples. Prediction of spontaneity of a reaction based on the cell emf. Numericals on standard electrode potential of half-cells, cell emf, relationship between free energy and equilibrium constant, standard electrode potential and free energy. (vi) Comparison of metallic conductance and electrolytic conductance. Relationship between conductance and resistance. Specific resistance and specific conductance. Cell constant: Calculation of cell constant. Meaning of equivalent conductance. Meaning of molar conductance. General relationship between specific conductance, molar conductance and equivalent conductance (units and graphs). Units, numericals. Molar conductance of a weak electrolyte at a given concentration and at infinite dilution. Kohlrausch’s Law – definition, applications and numericals. (vii) Faraday’s laws of Electrolysis. Faraday’s First Law of electrolysis. Statement, mathematical form. Simple problems. Faraday’s Second Law of electrolysis: Statement, mathematical form. Simple problems. Relation between Faraday, Avogadro’s number and charge on an electron. F = NAe should be given (no details of Millikan’s experiment are required). (viii) Batteries: Primary and Secondary Cells: Leclanche cell, mercury cell, Lead storage battery and fuel cell – structure, reactions and uses. (ix) Corrosion: Concept, mechanism of electrochemical reaction, factors affecting it and its prevention. Electrochemistry is the study of the processes involved in the inter conversion of electrical energy and chemical energy. Electrolysis is the process of decomposing an electrolyte in aqueous solution or in molten state by the passage of direct electric current. Positive ions (cations) migrate to the cathode. Negative ions (anions) migrate to the anode. Oxidation takes place at the anode and reduction takes place at the cathode. The reaction depends on the electron transfer at the electrodes and are called redox reactions. Faraday’s First During electrolysis, the amount of any substance deposited Law of or evolved at any electrode is proportional to the quantity Electrolysis: of electricity passed. W αQ Q=It, W= ZIt I - Current (ampere), t- time (sec), Q- amount of electricity passed, W- amount of substance liberated or deposited at the electrodes, z- electrochemical equivalent. Electrochemical Weight of the substance deposited or evolved by the Equivalent: passage of 1 coulomb of electricity. Faraday’s The weights of different substances evolved/ deposited by second Law of the passage of same quantity of electricity are proportional Electrolysis: to their chemical equivalents. W αE W1/E1 = W2/E2  Z1It/Z2It=E1/E2, E is proportional to Z ,Therefore E = FZ ,E= equal weight, F=Faraday (96500C), Z=Electrochemical eq. [ W= ItE/96500. This equation is used to find current strength or quantity of electricity. coulometer or Apparatus used to find current strength or quantity of voltameter. electricity. Relation F= NA. e between : NA= 6.023 x 1023, e = 1.6 x 10-19 Eg: Cu2++ 2e → Cu 2x 96500 : I mole (64.5 g) Q=nF , n= number of electrons gained, Galvanic cell (Voltaic cell) (Electrochemical cell): A device that produces an EMF as a result of chemical reactions that takes place within it which converts chemical energy into electrical energy. DIAGRAM: A zinc strip dipped in ZnSO4 solution and a copper strip dipped in CuSO4 solution are taken in two separate beakers.The electrodes are connected by the conducting wires through a voltmeter. At anode: oxidation Zn – 2e →Zn2+ At cathode: reduction Cu2+ + 2e →Cu Zinc being more electropositive than copper undergoes oxidation and acquires negative charge. Copper is less electropositive , undergoes reduction and it acquires positive charge. The elctrons flow from Zn rod to the Cu rod. However , the flow of electrons stop after sometime because the electrons in the Zn rod is attracted by the surrounding Zn2+ ions in the solution and the copper rod stops accepting electrons as it is surrounded by negatively charged sulphate ions. So to allow the flow of electrons , we use salt bridge. Salt bridge is a U tube filled with the solution of some inert electrolyte ( KCl,KNO3, or NH4Cl) to which gelatin or agar- agar has been added to convert into a semi solid state. The negative ions from the inert electrolyte will move towards the anodic compartment and neutralize the positive charge on it and vice versa. Functions of salt bridge: It allows the flow of current by completing the circuit. It maintains electrical neutrality. prevents mixing up of electrolytes. Representation of galvanic cell : Anode where oxidation takes place is written on the LHS. Cathode where reduction takes place is written on the RHS. Single The electrical potential difference set up between the electrode metal and the solution of its own ions is known as half potential cell electrode potential. It is the tendency of metal when placed in contact with its ions to become positively or negatively charged Oxidation Tendency of a metal to be negatively charged with potential respect to the solution. (tendency to lose electrons). M(s) → Mn+ + n e- Reduction Tendency of a metal to be positively charged with potential respect to the solution. (Tendency to gain electrons). Mn+ + n e- → M(s) Standard The electrode potential of an electrode determined in electrode relative to the standard hydrogen electrode under potential standard conditions. The standard conditions are IM concentration of ions at 298K (25°C) temperature and 1 atmospheric pressure. Measurement E° = E° cathode - E° anode of potential E0 cell = Reduction potential of cathode – reduction potential of anode. E° cell = E ° Right - E° left Electrode which has the higher value of reduction potential is the cathode. If E° cell is + ve , reaction is spontaneous. Electrical When a metal is placed in the solution of its own salt, double layer the chemical reaction takes place and finally a dynamic  equilibrium is established, because –ve or +ve charge developed on the metal attracts the +vely or –vely charged free ions in the solution. Due to this attraction, the +ve or –ve ions remain quite close to the metal. This a short layer of +ve or –ve ions is formed all around the metal. This layer is called Helmholtz electrical double layer. Standard SHE consists of a Platinum wire sealed into a glass hydrogen tube with a Pt foil at one end. The Pt foil is coated electrode/ with finely divided Pt. The electrode is placed in a normal beaker containing an aqueous solution of some acid hydrogen having an 1M concentrated of H+ ions. H2 gas is electrode. continuously bubbled through the solution. SHE is (SHE/ NHE) assigned ‘zero volt’ by convention. The metals above H2 undergoes oxidation with respect to H-cell and metals below H2 undergoes reduction with respect to H-cell.  Limitations It is not always convenient to use SHE as a reference electrode because : Unity concentration of H2 ions is difficult to be maintained,1 atmospheric pressure of H2 gas cannot be maintained uniformly, H2 – electrode gets poisoned even if traces of impurities are present. To overcome 1. Calomel electrode Hg – HgCl2 / KCl (aq) , Ag-AgCl / these KCl (aq) difficulties, 2. Weston electrode : -ve electrode Cd in CdSO4. +ve some other electrode Hg/ HgSO4 emf =1.0186v at 20 °C, very low electrodes temp. coefficient , easily reproducible. are used. Measurement of standard reduction potential/ standard electrode potential Reduction – Cathode Cell representation H+ + 1e→ H Zn / Zn2+ (aqueous) (IM) || H+ (IM) / H2 (1 H + H→ H2 atm) / Pt Anode -oxidation E° cell = E° H+/H2 - E°Zn2+ / Zn Zn – 2e→ Zn2+ = O - E° Zn2+ / Zn E° Zn2+/ Zn = - 0.76V To determine the cell representation : reduction potential of Pt / H2(1 atm) / H+ (IM) || Cu2+ (IM) / Cu. Cu(s) Cu – acts as the cathode. SHE – acts as the anode Anode : H2 – 2e → 2H+ Cathode: Cu 2+ + 2e→ Cu E° cell = E° Cu2+ / Cu - E° Pt/ H2 / H+ = 0.34 – 0 E° Cu2+ / Cu = 0.34V Factors affecting electrode potential : Nature of the The metals have different tendency to lose or electrode : accept the e -1 s and hence have different electrode potentials. The metals which are more electropositive has higher tendency to lose e-1s, thus it has low reduction potential but high value of oxidation potential. Concentration of ions Eg : Zn rod dipped in ZnSO4 solution as the concentration of the Zn2+ ions decreases electrode potential increases. Temperature Higher temperature makes the E° reduction more –ve. Nernst Equation. E red = Eored - 2.303 RT/nF log[reduced form]/[oxidised form] Electrode potential of R= 8.314 J/K /mol T=25 0C an electrode: F= 96500C n= no. of e gained. Ered = Eored - 0.0591/n log [M]/[Mn+] [M]= 1 Eg: Cu2++2e-> Cu E Cu2+/ Cu= EoCu2+/Cu- 0.0591/2 log 1/[Cu2+] Electrode Potential of Ecell= E0 cell- 0.0591/n log [oxidn]/[redn] a cell: Eg: Zn(S)/ Zn 2+ // Cu2+// Cu(s) Zn-2e→ Zn2+ Cu2++ 2e →Cu Zn+Cu2+->Zn2++ Cu Ecell= E0cell- 0.0591/2 log [Zn2+/Cu2+] Relationship Between At equilibrium, Ecell=0 Eocell and Equilibrium Ecell= E0 cell- 0.0591/n log k constant: Therefore E0 cell= 0.0591/n log k Relationship between Free energy (ΔG°) is the maximum energy E0cell and free available from the system which can be energy: converted into useful work. ΔG° = -nFE°cell If E0 cell is the positive, ΔG0 will be negative, the reaction is spontaneous Electrochemical series (Electromotive series): The arrangement of various elements in the order of increasing value of std. reduction potential. Important Characteristics: The most active metal are at the top of the series. Metals near the top of the series are strongly electropositive and lose e-1s readily to give cations i.e.. they are easily oxidised. Hence, they are good reducing agents. Greater the value of E0redn. Of a given species, greater its tendency to accept e-1s to undergo reduction and hence to act as oxidising agent. Eg: F2 is a strong oxidizing agent while Li+ is a weak oxidising agent. Metals above Hydrogen displace Hydrogen from dilute acids. To predict whether a given redox reaction is spontaneous or not. E0cell (+ve) , ΔG ( -ve) Greater the E0 reaction, lower its chemical reactivities of the metals. A metal with lower E˚redn will displace another metal/Hydrogen with higher std. redn potential gets precipitated. Conductors: The substances which allow the passage of electricity through them are called conductors. Metallic Conductors Electrolytes Movement of e-1 Movement of ions No change is observed in the Chemical change resulting in the chemical decomposition of electrolytes properties of the conductors Does not involve any transfer of Involves the transfer of matter as matter ions Resistance increases with increase Resistance decreases with increase in temp. in temp Follows Ohm’s Law Does not follow ohm’s law Factors affecting electrolytic conduction 1) Solute - solute interaction 2) Solute- solvent interaction 3) Viscosity of solvent 4) Temperature 5) Size of the ion 6) As temp. increases, the effect of the 1st 3 factors diminishes & K.E of the ions increases. This results in the increase of the electrical conductance of electrolytic solutions. Ohm’s law V∝I i.e V= I x R. V-Potential difference , I= Current, R= resistance of the conductor Resistance ‘R’ R∝l/a i.e R= P l/a It is the obstruction to the flow of current. P- resistivity, l- length of the conductor. a- area of cross section Resistivity is the resistance offered by the conductor of 1cm (specific length with the area of cross section equal to 1cm3 resistance) p= k a/l unit Ω cm Conductance is a measure of the ease with which current flows thru’ the conductor. C= 1/R unit Ω-1/mho/siemens(s) Specific is the reciprocal of specific resistance. Conductance (K=Kappa)K=1/p=1/R l/a unit :Ω-1 cm-1 (Conductivity or K is the conductivity of 1 cm cube of the soln. of an specific electrolyte conductivity / electrolytic conductivity Equivalent is the conductivity of a soln. containing 1gm.eq. of Conductivity the solute when placed between 2 electrodes 1cm apart. ^ eq = KxV , v= 1000/C ^ eq = K x 1000/C or ^ eq = K x 1000/N N is Normality unit : Ω-1cm2eq-1 Molar Conductivity of the solution containing 1 gm. Conductivity mole of the electrolyte such that the entire solution is placed between 2 parallel electrodes 1 cm apart. ^m = kx1000 /M ,M is molarity Unit = Ω-1 cm2 mol-1 Relationship ^ eq =^ m/z between ^ eq Z = no. of charges carried by the ions of the and ^m electrolyte. (valency of metal ion). Eg : ^ eq (MgSO4) = ^m (MgSO4)/2 Measurement of Cell constant is the ratio of the distance between 2 cell constant (l/a) electrodes and the area of the electrode. Unit – cm-1 Cell constant is obtained by measuring the conductance of a standard solution whose conductivity is known. Cell constant = Conductivity/Conductance Variation of specific conductance with dilution Specific conductance decreases as the no.of ions per unit area decreases. K depends on 1) no.of ions 2) amount of water present. As dilution increases, the conducting power of 1cm3 of the electrolytic solution decreases, owing to the fact there are lesser no. of ions and more H2O molecules present in 1cm3 of solution. Variation of eq. ^ eq or ^m increases as the no. of ions conductance with dilution present in a solution containing 1 g.eq or 1mole of electrolyte increases with dilution For strong electrolytes Conductance is very high and increases only slightly, but linearly with dilution  and soon reaches the maximum value, known as ^ eq or ^m at infinite dilution. Electrolyte is completely ionised in solution, and ^m∝ for strong electrolyte can be obtained by extrapolation of the curve to zero concentration. For weak electrolytes In concentrated solutions , the ^ eq or ^m is very low, and increases steadily with dilution. Here only a small fraction of molecules are in ionized form, and ^m∝ for weak electrolyte cannot be obtained by extrapolation of the curve to zero concentration. Degree of dissociation / ionization For strong electrolyte, KOHLRAUSCH’S LAW Debye- Huckel – Onsager At infinite dilution, when the dissociation of the electrolyte is complete, each ion equation makes a definite contribution towards the Λm c = Λm∝ - k √c molar conductivity of electrolyte, irrespective of the nature of the other ion with which it is associated.  where Λm c is molar This implies that the molar conductivity conductivity at a given of an electrolyte at infinite dilution can concentration be expressed as the sum of the Λm∝ is molar conductivity contributions from its individual ions. at infinite dilution. K is a If λo + and λo − represent the molar constant which depends conductivities of cation and anion upon the viscosity and respectively at infinite dilution, then the dielectric constant of the molar conductivity of electrolyte at solvent, C is the infinite dilution concentration of solution. Λm∝ = n+ λ α+ + n− λ α− where n+ and n− represent the number of positive and negative ions furnished by each formula unit of the electrolyte. Batteries It is the arrangement of two or more galvanic cells in series. Requirements of a good battery. 1. Light and compact. 2. Voltage should not vary during its use. 3.Should provide power for a longer time. 4. Should be rechargeable. Types of cell Primary Cells (or batteries) : Cell reaction is not reversible. In the primary batteries. the reaction occurs only once and after use over a period of time battery becomes dead and cannot be reused again. Secondary Cells : Cell reaction can be reversed by passing electricity through the cell (charging).  Flow Batteries and Fuel Cells : Materials pass through the battery which converts chemical energy to electric energy. Primary Uses Anode Cathode cell Leclanche used in Graphite rod cell watches , Zinc surrounded by radios and container MnO2 Cell calculators powder. potential Zn(s) → 1.25 V to Electrolyte-Paste of Zn2+(aq) + 2e- MnO2 + NH4+ + 1.5 V NH4Cl + ZnCl2 e → MnO(OH) + NH3 Mercury Used in Electrolyte:Moist paste Zn-Hg Paste of (HgO Cell the shape of KOH-ZnO amalgam + C) of button Cell cells for Zn (Hg) + 2OH- HgO + H2O +2e watches, → ZnO + H2O →Hg + 2OH- potential 1.35 V hearing aids, and calculator Secondary The cell cells reaction when the  battery is in use Lead Used in Electrolyte :Aq.soln of spongy lead grid of Pb storage automobil H2SO4 (38%) Pb + SO4 2 → packed with battery es and PbSO4 + 2e PbO2 invertors PbO2 + SO42- + 4H+ + 2e →PbSO4 Cell + 2H2O potential 2V per cell On charging the PbSO4 + 2e→ PbSO4 + 2H2O→ battery , reaction is Pb + SO4 2 PbO2 + SO42- + + 4H + 2e reversed and PbSO4 (s) on anode and cathode is converted into Pb and PbO2 respectively. Fuel Cell Used in porous C porous C rocket O2 (g) + 2H2O (l) Cell 2OH- [H2 (g) + + 4e- → 4OH- potential aq) → 2H2O (l) + 2e- ] x2 (aq) 0.6 V to 0.7 V Electrolyte : Conc KOH Lead storage battery :During the working of cell , conc of H2SO4 decreases as sulphate ions are consumed to form PbSO4.Anode and cathode are arranged alternatively.The cells are connected in series so as to increase the voltage of the battery. Fuel Cell : is an electrochemical cell that requires a continuous supply of reactants to keep functioning. Advantages of fuel cells Pollution free working, High efficiency (70-75%),Continuous source of energy.  Note:H2 –O2 Fuel cell has been used for generating power in Apollo space programme. ELECTROCHEMICAL THEORY OF RUSTING Corrosion :It is the process of slow conversion of metals into their undesirable compounds by reaction with moisture and other gases present in the atmosphere. Factors that affect corrosion (i) Reactivity of the metal (ii) Presence of carbon dioxide in water (iii)Air and moisture (iv)Strains in metals (v)Presence of electrolytes MECHANISM OF CORROSION Reactions at the anode and cathode Anode – pure iron ( iron in contact with water ) Cathode- Impure surface (iron in contact with air) Anode :Fe → Fe 2+ + 2e Cathode ; 2H + +2e → 2H 4Fe2+ +O2 +4H2O → 2Fe2O3 + 8H+ ; Fe2O3 +xH2O →Fe2O3.xH2O(rust) Prevention of 1.Barrier protection 2.Sacrificial protection corrosion 3.Electrical protection 4.Use of anti –rust solutions Barrier protection By painting the surface ,By coating the surface with a thin film of oil or grease ,By electroplating iron. Sacrificial protection Surface of iron is covered with a layer of more active metal like zinc. Zn loses electrons and prevents the rusting of iron.This process is galvanisation. Electrical protection The exposed surface of iron is protected by connecting it to more active metal like Mg. The more active metal acts as the anode and iron acts as the cathode Use of anti –rust The alkaline solutions of phosphate and chromate salts solutions. act as anti rust solutions. When iron articles are dipped in boiling and alkaline solutions of Na3(PO4), a protective insoluble film of iron phosphate is formed on them.