ELECTROCHEMICAL CELLS Structure 13.1 Introduction Objective8 13.2 Electrochemical or Galvanic Cells 13.2 1 A Convenient Notation for UIC Representation of El& and Ceb 13.3 Types of Electrodes 13.4 Electromotive Force and its Measurement 13.5 Free Energy Change and Electrical Work 13.6 The Nernst Equation 13.7 Standard Electrode Potentials 13.8 Using the Table of Standard Electrode Potentials 13.8.1 SigniT~canceof Positive and Negative Values 13.8.2 Elearode Potential and Stoichimetry 13.8.3 Displacement Reacriom 13.9 Calculation of the EMF of Galvanic Cells 13.10 Applications of EMF Measurements 13.1 1 Corrosion and its Prevention 13.12 Protective Measures Against Corrosion 13.13 Eec trochemical Energy Sources 13.3.1 Dry Cells 13.3.2 Nickel-CadmiumBatteries 13.3.3 Fuel Cells 13.14 Summary 13.15 Glossary 13.17 Answers to SAQs 13.1 INTRODUCTION Electrical energy can be used to bring about chemical reactions in electrolytic cells. The passage of electricity through aqueous solutions of electrolytes or molten electrolytes provides sufficient energy to cause an otherwise non-spontaneous reduction-oxidation or redox reaction to take place. This process, called electrolysis is widely used in electroplating, extractive metallurgy, electrochemical machining, etc. The converse process of conversion of the chemical energy of a spontaneous redox reaction into electrical energy takes palce in electrochemical or galvanic cells. These are also called voltaic cells in honour of Volta. He demostrated that by interposing sheets of paper soaked in salt water between different metals and connecting the metals, electricity could be produced. Batteries and fuel cells are such devices. The phenomenon of corrosion also belongs to this category. Both in electrolytic and galvanic cells, the redox reactions involve the transfer of charge between the electrode and the electrolyte. Thus, electrochemisuy can be broadly defined as the science that deals with the consequences of transfer of electric charges from one phase to another. Application of electrochemistry such as the extraction and refining of metals, manufacture of important chemicals like chlorine, caustic soda etc., are quite well known. The increasing awareness to minimise the pollution by automobile exhausts and thermal power plants has rekindled an intrest in batteries and fuel cells. Non-conventional energy sources like the solar energy or wind power can be used to generate electricity which can be used to charge the batteries. The batteries can be used during periods of break down of power and also in running automobiles and mopeds. 0 bjectives After studying this unit, you should be able to : * understand how a redox reaction can be made to produce electric current in a galvanic cell, * represent such a galvanic cell and to explain by means of electrode reactions how the cell produces an electric current, * identify the anode, cathode, psitive electrode and negative electrode of a' galvanic cell, * know how the standard electrode potentials are determined and how to use these to calculate the standard E.M.F. of a cell, * know how to use the relationship between free energy change and E.M.F. and to calculate the equilibrium constant from the standard E.M.F. of a cell, * use Nernst equation to calculate the electrode potential of half cells or the E.M.F. of cells in which the reactants and products are not in their standard states. * Explain the working of a pH meter and use EMF measurements for measuring pH and solubility products, * appreciate the usefulness of potentiometric titrations, * Analyse the advantages and limitations of various electrochemical energy sources, and * explain what galvanic corrosion is and how it may be minimized. These are devices in which the free energy change accompanying a redox reaction is converted into electrical energy. If a zinc rod is dipped into a solution of copper sulphate, a brownish-red deposit of copjgr is formed on the surface of zinc and the bluish-green colour of the copper sulphate solution disappears. The redox reactions that occur can be represented as oxidation : ~n(s)-2e+~x?+(aq) (13.1) reduction : cu2+(aq)+2e+Cu(s) (13.2) The overall reaction is represented as (13.3) Zn(s) +C$+(aq) +~x?+(aq) +Cu(s) (13.3) In reaction (13.3), only the active constituents taking part in the reaction are indicated. The sulphate ions do not take part in the reaction. At 298 K and 101.3 kPa (i.e., lam) the standard free energy change ( AGO ) corresponding to this spontaneous reaction is -212.9 kJ. Though AGO is a measure of the electrical work that can be derived from this reaction i.e., Eq.(13.3), the electron transfer from zinc to Cu2+( aq ) can be demonstrated only when the oxidation and reduction reaction are allowed to take place separately as in the Daniel1 cell, which is a Galvanic cell. [Figure 13.11. This cell consists of a zinc rod (electrode) dipping into a solution of zinc sulphate (one can use an aqueous solution of sodium chloride also) and a copper rod (a platinum electrode can also be used ) dipping into a solution Of capper sulphate. A porous barrier prevents the mixing of the two electrolytes [ Figurel3.l(a)] but allows the passage of the ions. The two electrolytes may also be kept in separate beakers and a salt bridge enables the passage of ions from one compartment to another [Figure 13.1(b)]. When the two electrodes are connected by means ~f a wire (electronic conductor), oxidation (reaction-1) takes palce at the zinc electrode. The zinc ions dissolve in the Elwetron flow EMr~&emlerlCdls Porous borrlrr !(a) Solutlow equated by a porous bad= so$- - (b) Sdutlons separated by a salt brldge Flpre 13.1 :Laboratory Verdons of the Daniell Cell electrolyte and the electrons left behind on the electrode push other electrons via the connecting wire to the copper electrode. These electrons are used by cu2+ions in solution and reduction (reaction-2) occurs. The electron flow from the zinc electrode to the copper electrode provides a source of electricity. The direction of the current flow, as measured by a current measuring device, is opposite to that of electron flow. Sirice oxidation takes place at the zinc electrode it is called the anode. In this galvanic cell, this electrode, being a source of a negative charge, is the negative electrode. The copper electrode at which reduction occurs is the cathode and since it is a sink for electrons (accepts electrons), it is the positive electrode. As the cell continues to produce current, the zn2+ions entering the electrically neutral solution make the solution positively charged. The zinc electrode acquires a nagative charge with respect to the solution. As a result of the build up of charges, an electrical double layer is established at the electrode-electrolyte interface. The resultant potential difference between the electrode and electrolyte is called the electrode potential. The negative electrode potential on the zinc electrode will prevent the zinc ions from the metal lattice leaving the electrode. At the other electrode, the cuZ+ions in solution are used up in the reduction (reaction-2). The remaining sulphate ions tend to make the solution Acquire a negative charge. Since electrons are removed from the copper electrode it acquires a positive potential and so prevents the approach of cu2+ions to the electrode (cathode). A continuous flow of current can be maintained if electrical neutrality is maintained at both the compartments A and C. The excess of sulphate ions from C move through the barrier into A so as to neutralize the excess positive charge of the solution in A. The sulphate ions do not undergo any chemical change. Instead of a porous partition, a salt bridge is often used. This consists of a glass tube containing a concentrated solution of KC1 or NH,NO, gelled by adding gelatin or Equilibria & Electroehmlstry agar. The gel confines the electrolyte to the tube, prevents the mixing of the electrolytes and allows the passage of ions. Here, K+or Neions move to C and the anions ( Cl- or NO; ) to A so as to maintain electrical neutrality. The high concentration ensures that most of the current is carried by the ions of the electrolyte used in the salt bridge. These ions also carry equal shares of the current and do not undergo any reaction with the electrode. 13.2.1 A Convenient Notation for the Representation of Electrodes and Cells Any electrochemical cell like the Daniell cell consists of two electrodeelectrolyte assemblies and each of these is called a half cell or simply an electrode. Oxidation occurs at one half cell and reduction at the other. The cell reaction is simply the sum of the two half cell or electrode reactions. A convenient notation is generally followed to represent on paper the electrodes and cells. A single vertical line denotes the boundary between two phases. In the case of aqueous solutions, the concentrations of the ionic species are indicated in parenthesis. The Daniell cell [Fig.13.1] can be represented on paper as ~n 1 w+(1.0~1 II cu2+( 1.0~) I cu The double vertical line between the electrolyte solutions indicates that a barrier [Figure 13.l(a)] or a salt bridge [Figure 13.1 (b)] has been used. The half cell or electrode at which oxidation occurs, i.e., anode, is written on the left hand side, and the cathode on the right hand side. An easy way of remembering all that has been said about galvanic cells is given below, using the Daniell cell as an example. L H S electrode B H S electrode Qxidation occurs Reduction occurs Anode Cathode Uega tive Electrode Positive Electrode It is simpler to remember the letters "LOAN" which summarizes the conventions.