11Chapter 21aElectrochemistry: The Electrolytic Cell2ElectrochemistryElectrochemical reactions are oxidation-reduction reactions.| The two parts of the reaction are physically separated.z The oxidation reaction occurs in one cell.z The reduction reaction occurs in the other cell.3Electrochemistry| There are two kinds electrochemical cells.1. Electrochemical cells containing spontaneous chemical reactions are called voltaicorgalvanic cells.  The generation of electric current from a chemical reaction.2. Electrochemical cells containing in nonspontaneous chemical reactions are called electrolytic cells. The use of electric current to produce a chemical change.4Electrical Conduction| Metals conduct electric currents well in a process called metallic conduction.z In metallic conduction there is electron flow with no atomic motion.z Metal atoms changing oxidation states without moving.• E.g. Oxidative phosphorylation5Electrical Conduction| In ionic or electrolytic conduction ionic motion transports the electrons.z Positively charged ions, cations, move toward the negative electrode.z Negatively charged ions, anions, move toward the positive electrode. 6ElectrodesThe following convention for electrodes is correct for either electrolytic or voltaic cells:| The cathode is the electrode at which reductionoccurs.• The cathode is negative in electrolytic cells and positive in voltaic cells.| The anode is the electrode at which oxidationoccurs.• The anode is positive in electrolytic cells and negative in voltaic cells.27Electrodes| Inert electrodes do not react with the liquids or products of the electrochemical reaction.| Two examples of common inert electrodes are graphite and platinum.8Electrolytic CellsElectrical energy is used to force nonspontaneous chemical reactions to occur.| The process is called electrolysis.| Two examples of commercial electrolytic reactions are:1. The electroplating of jewelry and auto parts.2. The electrolysis of chemical compounds.9Electrolytic Cells| Electrolytic cells consist of:1. A container for the reaction mixture.2. Two electrodes immersed in the reaction mixture.3. A source of direct current.| Electrolytic cells uses electrical energy to produce a chemical change.z The electrical energy forces a current through a cell that has anegative potential.z The electrical energy forces a chemical change to occur.10Figure 11.19:(a) A standard galvanic cell (b) A standard electrolytic cellThe cell in (b) has a power source that forces the electrons in the opposite direction from the voltaic cell in (a).Electrolytic CellVoltaic Cell11Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| The stoichiometry of electrolysis processes can quantify “how much chemical change occurs with the flow of a given current for a specific time”.12Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| Faraday’s Law - The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell.| A faraday is the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.23 -1 6.022 10 faraday of electricity e≡×313Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| A coulomb is the amount of charge that passes a given point when a current of one ampere (A) flows for one second.| 1 ampere (amp) = 1 coulomb/second23 -––1 6.022 10 1 1.01.0 96, 485 faraday efaraday mol emol e coulombs≡×≡≡14Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| Faraday’s Law states that during electrolysis, one faraday of electricity (96,485 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent.z This corresponds to the passage of one mole of electrons through the electrolytic cell.23 –23 –1 6.022 10 1 6.022 10 equivalent of oxidizing agent gain of eequivalent of reducing agent loss of e≡×≡×15Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| Example: Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes. ()2- 0: 2 1 2 1 106 2(96, 485) 106 3.20 3.20 1 96,48560 3.20 1? 30.0 minmin 96, 485Cathode Pd e Pdmol mol molggCampsmol e CsCmolegsC+−−+→== =  -106 3.16 2mol Pd g PdgPdmol e mol Pd=16The Electrolysis of Water| Hydrogen and oxygen combine spontaneously to form water.z The decrease in free energy that accompanies this spontaneous reaction can be used to run fuel cells to produce electricity.| The reverse process, which is not spontaneous, requires energy to occur.| The formation of oxygen and hydrogen gases from water can be forced by electrolysis.17The Electrolysis of Water222()22()22()2()4 22()2( 2 4 4 2(2 2 2 ) 6 2 4 4 2 2 ggggHOggAnode reaction H O O H eCathode reaction H O e H OHCell reaction H O H O H OHThe overall reaction is H O H O++→+ ++→ +→+++→+––––) 18Counting Electrons: Coulometry and Faraday’s Law of Electrolysis| Example: Calculate the volume of oxygen (measured at STP) produced by the oxidation of water during the passage of 3.20 amperes of current for 30.0 minutes. ()()()-222: 2 4 4 2 1 4 4 22.4 4 96,500 ? 3.20 3.20 196,4851.0 22.460 30.0 minmingSTPSTPSTPAnode H O O H emol mol mol molLCCLO ampsmol e Cmol Ls+−→++====22-22222.4 3.20 196,485 4 0.334 334 STPSTP STPLOCmole molOsCmolemolOLOor mL O− =  =419The Electrolysis of Molten Sodium Chloride| Liquid sodium is produced at one electrode.z Indicates that the reaction Na+(A)+ e-→ Na(A)occurs at this electrode.z Is this electrode the anode or cathode?z Reduction occurs at the cathode.| Gaseous chlorine is produced at the other electrode.z Indicates that the reaction 2 Cl-→ Cl2(g)+ 2 e-occurs at this electrode.z Is this electrode the anode or cathode?z Oxidation occurs at the anode.20The Electrolysis of Molten Sodium Chloride| In all electrolytic cells, electrons are forced to flow