Electrical Conductance Substances that allow the passage of current through them are called conductors and the phenomenon is called electrical conductance. Metallic: the conduction is only due to the presence of free mobile electrons. Electrolytic: The substances which conduct electricity both in the fused state and in the aqueous solution are called electrolytic conductors or electrolytes. Types of Electrolytes Based on the extent of dissociation, electrolytes are classified into two types: Strong electrolytes: For example, HCl, HNO3, NaOH, KOH, NaCl and KCl. Weak electrolytes: For example, CH3COOH, HCOOH and NH4OH.

Mechanism involved in typical electrolytic conduction. Electrolytic Conductance Mechanism involved in typical electrolytic conduction.

Specific Conductance Electrolytic conductor offers some resistance to the flow of current. Let R be the resistance offered by an electrolytic conductor. This resistance is directly proportional to the length (l) of the conductor and inversely proportional to the area of cross-section (a) of the conductor. R∝ l and R∝ 1 a So, R∝ l or R = ρ. l , where ρ is the proportionality constant. a a Therefore, ρ = R. a/l If R= 1Ω, l = 1m and a = 1 m2, then ρ = R, where ρ is the specific resistance. It has unit of Ω m (ohm- meter)in SI system.

The reciprocal of specific resistance is called the specific conductance (k). Specific conductance is the conductance of the solution placed between two electrodes having the cross-sectional area of 1 m2 and placed 1 m apar.t. In SI system, it has unit of Siemens per centimeter (S/cm) The ratio of length to the cross-sectional area is called the cell constant

Equivalent Conductance – It is the conductance of an electrolytic solution due to all the ions obtained from 1 gram-equivalent mass of the electrolyte at a given concentration. SI Unit = Sm2/equiv. Molar Conductance SI Unit = Sm2/mol

Q1. Calculate the specific conductance (k) of a solution which is placed between two Electrodes 1.63 x 10-2 m apart and having cross-sectional area 3.9996 x 10-4 m2. The resistance offered by the solution is 17.69 ohm. Ans. = 2.3058 S/m Q2. At 298 K, the specific conductance of 0.1 N NaCl is 1.1 S/m. Calculate the equivalent conductance at the same temp. Ans. = 1.1 x 10-2 Sm2/equiv. Q3. Calculate the specific conductance and molar conductance of 0.0469 M NaOH, if the resistance offered and the cell constant are 60.2 ohm and 89.6 m-1, respectively. Ans. = 0.03173 Sm2/mol

Cells The devices which convert electrical energy into chemical energy or vice versa are called cells. Based on the activity taking place in them, these devices are classified into two major categories: Electrolytic cells. Galvanic (voltaic cells): These are further divided into chemical cells; concentration cells. The devices in which chemical changes occur in the presence of applied electrical energy are referred to as electrolytic cells. The devices in which electrical energy is generated on account of the chemical reactions occurring in them are known as galvanic cells.

Redox Reactions Such reactions in which oxidation and reduction take place simultaneously are known as redox reactions. For example, in the reaction:

Electrode Potential Origin of Electrode Potential Origin of electrode potential (reduction) Origin of electrode potential (oxidation)

Standard Electrode Potential (E0) Oxidation Potential If oxidation occurs at the electrode, at equilibrium the potential of the electrode can be termed as oxidation potential. For the zinc electrode, the reaction can be represented as: Reduction Potential If reduction occurs at the electrode, at equilibrium the potential of the electrode is referred to as reduction potential. For the zinc electrode, the reaction can be represented as: Standard Electrode Potential (E0) If is defined as the potential that exists between the metal or the gas and its aqueous solution of unit concentration at 298 K when the sum of all partial pressures of the gaseous reactants and products, if any, is equal to 1 atm pressure.

Galvanic Cells The cell consists of two half-cells. Oxidation takes place in one half-cell and reduction in the other as indicated by the half-reactions

Salt Bridge

EMF of the Cell and Free Energy Change A salt bridge helps to bring about internal contact between the electrodes; to minimize liquid junction potential; to minimize polarization. EMF of the Cell and Free Energy Change The difference in potential of the electrodes is known as electromotive force.

Electrochemical Conventions and Notations There are two electrodes in a galvanic cell. Each of these electrodes is referred to as a half cell. If reduction occurs at an electrode, it is called a reduction electrode or a positive electrode. If oxidation occurs at an electrode, it is called an oxidation electrode or a negative electrode. Oxidation and reduction potential values of an electrode are numerically the same, with opposite signs. As per IUPAC, the potential of an electrode is always expressed as reduction potential only, irrespective of the reactions occurring at the electrode.

The interphase across which a potential develops is denoted by either a single vertical line (|) or a semicolon (;). Representation of a cell will have the oxidation electrode (anode) on the LHS and the reduction electrode (cathode) along the RHS. The EMF of a cell is represented as:

Nernst equation for an electrode The potential of an electrode depends on concentration and temperature. A quantitative equation relating the electrode potential with these parameters was deduced by Walter Herman Nernst. Nernst equation for an electrode Nernst equation for a galvanic cell At Anode M1 → M1n+ + ne- At Cathode M2n+ + ne- → M2 M1 + M2n+ → M1n+ + M2

Q1. For the cell Zn I Zn2+ (a= 10-4M) II Mg2+(a=10-3 M) I Mg, the standard reduction potential for Zn and Mg electrodes are -0.764 V and -2.364 V, respectively. Write the half cell reaction and overall cell reaction Find Eocell and  G and predict if the cell reaction is spontaneous or not. Q2. Calculate the standard electrode potential of lead electrode. If the electrode potential is -0.18025 V at 301K and a concentration of Pb2+ solution is 0.096M.

Q. Calculate the e.m.f of the cell in which the following reaction takes place: Ni(s) + 2Ag+ (0.002 M) → Ni2+ (0.160 M) + 2Ag(s) Given EoCell = 1.05 V

Application of Nernst equation The potential of an electrode and EMF of a cell can be calculated at any temperature and concentration. Knowing potential of an electrode, the concentration of the reactant can be calculated. The concentration of a solution in the galvanic cell can be determined. The pH of a solution can be calculated by measuring the EMF.

Thermodynamics of an Electrochemical Process Electrical Energy in a Galvanic Cell- The electrical energy produced by a galvanic cell is given by the product of its electromotive force and the quantity of electricity which passes. If EMF is measured in volts and quantity of electricity in coulombs, the electrical energy is obtained in Volts-coulombs or joules. For example in Daniell cell EMF is 1.10 volts and the cell reaction involves libration as well as taking up of 2 electrons therefore the quantity of electricity produced according to Faraday’s second law is 2 faradays i.e. 2 x 96500 coulombs. Hence the electrical energy generated in the cell for the complete cell reaction= 2 x 96500 x 1.1 = 212300 volt- coulombs i.e. joules The relation between EMF of the cell and the thermodynamic parameters H and S can be derived based on this fundamental relation between G and EMF:

Electrical Energy and Free Energy Change of the cell Reaction- According to Gibbs and Helmholtz, the electrical energy of a reversible cell is given by the free energy decrease( - G) of the reaction occurring in the cell. Suppose in a perticular cell reaction, n is the number of electrons liberated at one electrode (or taken up at the other electrode). Then evidently, n faradays (nF) of electricity will be generated in the complete cell reaction. Ecell is the EMF of the cell, then, Electric energy produced by the cell = nFE Hence,

Q For the Daniell cell involving the cell reaction Zn(s) + Cu2+(aq) Zn2+ (aq) + Cu(s) The standard free energies of formation of Zn, Cu, Cu2+ and Zn2+ are 0, 0, 64.4 Kjmol-1 and -154.0 kJ mol-1, respectively. Calculate the standard EMF of the cell. Standard EMF and Equilibrium Constant - Go = RT ln K = nFEo

Measurement of EMF of the Cell EMF and Potential Difference The terms EMF and potential difference are used indiscriminately, but there is a clear distinction between them. Differences between EMF and potential difference

Standard Electrode Potential (E0) It is defined as the potential that exists between the metal Or the gas and its aqueous solution of unit concentration at 298K when the sum of all partial pressures of the gaseous reactants and products, if any is equal to 1 atm pressure.

Standard Hydrogen Electrode (Normal Hydrogen Electrode) If reduction occurs at the electrode, the reaction taking place will be: The half-reaction is

As the potential cannot be measured directly, `it is conventionally assumed to be zero at 298 k. The EMF of SHE is assumed to be zero, and the potential of the second electrode can be obtained by constructing a cell with SHE. The potential of an electrode measured at standard condition is known as standard electrode potential. It may be mentioned, however that it is not convenient to use the standard hydrogen electrode as the reference electrode. This is because it is difficult to maintain the activity of H+ ions in the solution at unity and to keep the presence of the gas uniformly at one atmosphere. A far better reference electrode is the calomel electrode.

However, SHE/NHE has its own inherent limitations: Construction of SHE is a difficult activity. It is difficult to maintain unit molar concentration of hydrogen throughout and to pass hydrogen always at exactly 1 atm pressure. Presence of arsenic compounds would easily get absorbed on platinum foil, thereby poisoning the surface. It would affect equilibrium of the reaction. In addition, SHE cannot be used in presence of strong oxidizing and reducing agents.

Case 1: If the experimental electrode is an oxidation electrode, the cell so constructed shall be E1 = ESHE – Eexp. ESHE = 0 E1 = 0 - Eexp E exp = - E1

Case 2: If the experimental electrode is a reduction electrode, then the cell formed is represented as:

Characteristics of Electrochemical Series When the electrodes are arranged in the increasing order of their standard reduction potential values, the series so formed is termed as electrochemical series. Characteristics of Electrochemical Series Lithium is the first member of the series. Highly reactive metal systems are at the top of the series. In other words, good reducing agents are at the top of the series. All good oxidizing agents are at the bottom of the series. Hydrogen system is at the middle of the series.

Applications of Electrochemical Series Higher the reduction potential, greater is the tendency of the element to get reduced. Lower the reduction potential, greater is the reducing ability. A metal placed higher in the series is anodic to other metals which lie below it. Knowledge of electrochemical series helps in selection of electrode assemblies, to construct the galvanic cells of the desired EMF. The polarity of the electrode system and the electrode reaction can be easily predicted. Spontaneity and feasibility of the cell under construction can be easily predicted.

Electrochemical series Li K Ca Na Mg Al Zn Fe Ni H2 Cu Ag Pt Au F- Increasing order of standard reduction Potential (E0)

Characteristic of Electrochemical Series- The Eo values apply to the half cell reactions as read in the forward (left to right ) direction. The more positive Eo defines the greater tendency for the substance to be reduced e.g. F2(1 atm) + 2 e- 2F- (1M) Eo= 2.87 V F2 is the strongest oxidizing agent because it has the greatest tendency to be reduced while F- is the weakest reducing agent Similarly, the more negative value of Eo defines the less tendency to be reduced e.g. Li+ (1 atm) + 2 e- Li(s) Eo= -3.05 V thus Li+ is the weakest oxidizing agent because it is most difficult species to reduce and Li metal is the strongest reducing agent. Under standard state conditions, the oxidizing agents (the species on the left side of the half cell reactions) decreases in the strength from bottom to top and the reducing agents ( the species on the right hand side of the half-reacation) increase in strength from top to bottom.

Therefore Zn spontaneously reduces Cu++ to form Zn++ and Cu Under standard –state conditions any species on the left of a given half-cell reaction will react spontaneously with a species that appears on the right of any half-cell reaction located below it in Electrochemical series. This is sometimes called diagonal rule. Therefore Zn spontaneously reduces Cu++ to form Zn++ and Cu 4. Changing the stoichometric cofficients of a half cell reaction does not affect the value of Eo because electrode potentials are intensive properties. This means that the value of Eo is unaffected by the size of the electrodes or the amount of solutions present e.g. I2 (s) + 2e- 2I- (aq) Eo= 0.53 V Eo does not change if we multiply the half cell reaction by 2 2I2 (s) + 4e- 4I-(aq) Eo= 0.53 V 5. Like ∆ H, ∆ G and ∆S, the sign of Eo changes but magnitude remains same when we reverse a reaction. Zn2+ (aq) + 2e- Zn(s) Cu2+ (aq) + 2e- Cu(s)

Reversible electrochemical cells are the cells whose cell reactions can be get reversed when an external emf greater than its capacity is applied. (A cell which obeys thermodynamic conditions of reversibility is known as reversible cells). For example Daniel cell is a galvanic cell with capacity of 1.1 V, when an external emf of 1.1 V is applied, the cell reaction stops. Zn + Cu2+ Zn2+ + Cu But when an increased amount of emf greater than 1.1 V is applied, the cell reaction is get reversed. Zn2+ + Cu Zn + Cu2+ Irreversible cell A cell is irreversible if the cell reaction cannot be reversed e.g., the cell Zn H2SO4 Cu

Concentration Cells There is, however, another category of cells in which the EMF arises not due to any chemical reaction but due to transfer of matter from one half- cell to the other because of a difference in the concentrations of the species involved. These are called concentration cells. These are of two types Electrode concentration cell Electrolyte-concentration cell Electrode Concentration Cells In these cells, the concentration of the electrolyte is the same. The two electrodes contains the same substance but with different concentration.

Electrolyte- Concentration Cell In these cells, the electrodes contain different concentration of the same electrolyte. The two electrodes are connected directly through a diffusion membrane. Zn(s), Zn2+ (a+)1 II Zn2+(a+)2, Zn(s) RHE Zn2+(a+)2 + 2e- Zn(s) LHE Zn(s) Zn(a+)1 + 2e-