Electrolytic CellsVoltaic cells or Galvanic cells are driven by a spontaneous chemical (redox) reaction that produces an electric current through an external circuit. Although these voltaic cells are important, they are not the only type of electrochemical cell. An electrolytic cell drives an electric current through the system while doing work on the chemical system itself. Electrolysis is the process within the electrolytic cell that drives an oxidation-reduction reaction in the opposite direction of the Voltaic cell and therefore is non-spontaneous.Electrochemical vs. Electrolytic cellsSpontaneous vs. non spontaneousElectrochemical cells are powered by redox reactions which are spontaneous. These spontaneous redox reactions produce electrical energy that is harnessed in a battery. The reverse reaction in each case is non spontaneous and requires electrical energy to occur. The general form of the reaction can be written as:Spontaneous ---------->ReactantsProducts+Electrical Energy<----------- Non spontaneousWorking of an Electrolytic cellTo explain what happens in an electrolytic cell let us examine the electrolysis of molten sodium chloride into sodium metal and chlorine gas. The reaction is written below.---------> Non spontaneous ( electrolytic cell )2 Na Cl (l)2 Na (s)+Cl2 (g)<--------- Spontaneous ( electrochemical cell )If molten NaCl(l) is taken in a container and two inert electrodes of C(s)(carbon rods) are inserted, which are attached to the + and - terminals of a battery, an electrolytic reaction will occur. Examine the diagram below of the cell.Electrolysis takes place in an electrolytic cell, the simplest form of which is shown here on the right:The components which make contact with the electrolyte are called ELECTRODES. The electrode which is attached to the negative pole of the battery, and which supplies electrons to the electrolyte, is called the CATHODE. Reduction takes place at the cathode.The electrode which is attached to the positive pole of the battery, and which accepts electrons from the electrolyte, is called the ANODE. Oxidation takes place at the anode.When a direct electric current is passed through an ELECTROLYTE (such as a molten salt or an aqueous solution of a salt, acid or base), chemical reactions take place at the contacts between the circuit and the solution. This process is called ELECTROLYSIS. The combination of container, electrodes and electrolyte constitute the ELECTROLYTIC CELL. In such a cell, electrical energy from the current supply (battery or DC generator) is converted to chemical energy.Various reactions take place at the electrodes during electrolysis. In general, reduction takes place at the cathode, and oxidation takes place at the anode.Electrolysis is a hugely important process - many elements were discovered by electrolysis of their salts (the alkali metals, for example), and it is used industrially in a number of ways:The extraction of metals, notably aluminum, magnesium, and sodium.The preparation of halogens, for example chlorine.The refining of metals, such as copper and zinc.At the electrodes, ions become discharged, or the anode may become oxidized and pass into solution. INERT ELECTRODES are electrodes which do not undergo any change during electrolysis. Platinum and carbon are frequently used when inert electrodes are required. Electrolysis of Molten Sodium ChlorideWhen molten sodium chloride is electrolyzed the products obtained are, sodium metal at cathode and chlorine gas at anode.The electrolysis of molten sodium chloride is as follows:These ions carry the current, Na+ ions move towards cathode, and Cl- ions move towards anode. On reaching the respective electrodes these ions get involved in electrode reactions as follows. Electrolysis of Aqueous Sodium Chloride SolutionWhen sodium chloride solution is electrolyzed using platinum electrodes, the products are hydrogen gas at the cathode, and chlorine gas at anode. The mechanism of electrolysis involves the following reactions.Thus, Na+ and Cl- ions carry the current. Na+ ions go towards cathode and Cl- ions move towards anode.At anode: Cl- ions get oxidised to give chlorine gas.Cathode materialFavoured cathode reactionPlatinumH+ + e - ------> 1/2 H2(g)MercuryNa + e - ------> NaNa+ ions move towards cathode, but do not get reduced (due to its highly negative electrode potential). Instead, H+ ions produced due to the self-ionization of water get reduced at cathode.Na+ ions migrated to cathode, produce Na+OH- with OH- liberated at cathode.It is important to note that the products of electrolysis also depend upon the material used in making the electrodes. For example, during the electrolysis of aqueous solution of sodium chloride shown earlier, the products obtained are different when different electrodes are used.Electrolysis of Copper Sulphate SolutionWhen a solution of copper sulphate is electrolyzed using platinum electrodes, the products of electrolysis are copper at cathode and oxygen gas at anode. After electrolysis the solution around the anode is found to contain sulphuric acid. This can be explained as follows:The electrolyte CuSO4 when dissolved in water dissociates to give Cu2+ and SO42-ions.Current is carried by Cu2+ and SO42-in the solution. Cu2+ ions go to cathode and get reduced to metal copper.The SO42-ions move towards anode but is very difficult to get oxidized. So, at anode, oxidation of OH-ions (produced by the self-ionization of water) takes place in preference to SO42-. So, the anode reaction is,Thus, during the electrolysis of CuSO4 solution, although the current is carried by Cu2+ and SO42-ions, but the ions, which are involved in the electrode reactions are Cu2+ and OH-. Liberation of H+ around the anode makes the solution around it more acidic.Let us consider what happens when a fairly concentrated solution of copper(II) chloride (CuCl2) is electrolyzed in the cell shown below, where the electrodes are made of carbon:When the current is flowing, the following is seen to happen:Metallic copper is deposited on the cathodeChlorine gas is evolved at the anode.At the cathode, electrons are supplied to cations, which migrate to the cathode. The Cu2+ cation is discharged by accepting electrons. At the anode, electrons are supplied to the anions, which migrate to the anode. The half reactions areGenerally speaking, prediction of the nature of the products is complicated by (i) concentration effects and (ii) the rate at which the oxidation and reduction of different ions takes place. Standard electrode potentials can give a rough guide as to what may happen.As a guide, in fairly concentrated aqueous solution, metals with a positive (reduction potential) will be formed at the cathode. Otherwise, hydrogen is formed by reduction of water or of the H+ ion. Halogens (chlorine, bromine and iodine) are formed at the anode when aqueous solutions of halides are electrolyzed. The sulphate ion is not discharged (oxidized) at the anode in aqueous solution, rather, oxygen is formed by oxidation of water or of the OH- ion. Thus, electrolysis of sulphates produce oxygen at the anode.Electrolysis of water:Water may be electrolyzed in the apparatus shown below. Pure water is however a very poor conductor of electricity and one has to add dilute sulphuric acid in order to have a significant current flow.The electrodes consist of platinum foil. The electrolyte is dilute sulphuric acid. Hydrogen gas is evolved at the cathode, and oxygen at the anode.The ratio, by volume, of hydrogen to oxygen, is exactly 2:1.Remember that electron flow in the circuit is opposite to the conventional current flow. Thus, while the conventional current flows from the positive pole through the electrolyte and ends up at the negative pole, electrons flow from the negative pole in the reverse direction. The positive pole of a battery accept electrons from the electrolyte by means of the anode of the electrolytic cell. The reaction at the cathode (tube A) is the reduction of protons:Oxidation takes place at the anode (tube B). There are two anions competing to give up their electrons, namely sulphate (SO42-), and hydroxide (OH-) from the ionization of water. The oxidation of OH- according to the reactionhas a standard electrode potential of -0.40V, compared to the oxidation of sulphate (-0.17V), and consequently, OH- will be oxidized preferentially. The overall reaction is thereforeand the electrons are reurned to the battery, thus completing the circuit.Extraction of Aluminum:Aluminium is obtained by the electrolytic reduction of its molten oxide, alumina (Al2O3). Because alumina has a very high melting point (2045 ºC), the mineral cryolite (Na3AlF6) is added to lower the melting point in order that the electrolysis may be carried out at about 950 ºC. The electrolytic cell has carbon anodes and a carbon cathode (which forms the lining of the tank in which the electrolysis takes place). Carbon dioxide is formed at the anodes, and aluminum at the cathode. It is heavier than the molten alumina/cryolite mixture, and sinks to the bottom of the cell, where it is tapped off. The procedure is known as the Hall-Héroult process.Aluminium extraction is very demanding on electrical current (typically, 3-5 V and 100 000 A), and is economical only where power is cheap.Electro refining of copper:When copper is first obtained by reduction of its ores, it is cast as impure slabs or ingots, called blister copper. In the electro refining process, the blister ingots are used as anodes in an electrolytic cell, where an acid solution of copper (II) sulphate is used as electrolyte. Initially, the cathodes consist of thin sheets of pure copper.During electrolysis, copper passes into solution from the anodes, (leaving the impurities, normally containing silver, gold and platinum) as ananode slime, which sinks to the bottom of the cell. The anode reaction isAt the cathode, copper (II) ions are discharged and the pure copper sheet becomes coated with an increasingly thick layer of very pure copper:Electroplating:Electroplating consists of depositing a thin layer of a metal on another, either for protection or for the sake of appearance. Typically, a brass or nickel object is coated with a layer of silver by making use of electrolysis of a silver solution, using the object to be coated as the cathode:The anode consists of pure silver, and the cathode is the object to be plated. The electrolyte is a mixture of silver nitrate with potassium cyanide.The reactions are:At the anode: Ag → Ag+ + e-At the cathode: Ag+ + e- → AgThe cyanide ensures a low concentration of silver ions, a condition for providing the best plating results.During the process, the concentration of silver in the electrolyte remains constant, as the rate of reduction at the cathode (which is the rate of deposition of silver on the object) is the same as the rate of reduction at the anode (which is the rate of rate of dissolution of the silver anode). Electrochemical cellsElectrolytic cellsCATHODEAccepts electrons from the external circuit (wire).Accepts electrons from the external circuit (wire).Is the site of REDUCTION.Is the site of REDUCTION.Is the conventional POSITIVE pole of the cell.Is attached to the NEGATIVE pole of the DC source.ANODEReleases electrons to the external circuit (wire).Releases electrons to the external circuit (wire).Is the site of OXIDATION.Is the site of OXIDATIONIs the conventional NEGATIVE pole of the cell.Is attached to the POSITIVE pole of the DC source.