Slide 1Slide 2Oxidation numberSlide 4Slide 5Slide 6Slide 7Slide 8Slide 9Slide 10Slide 11Slide 12Slide 13Slide 14Slide 15Slide 16Slide 17Slide 18Slide 19Slide 20Slide 21Slide 22Slide 23Slide 24Slide 25Slide 26Slide 27Slide 28Slide 29Slide 30Slide 31Slide 32Slide 33ElectrochemistryChapter 18Copyright©TheMcGraw-HillCompanies,Inc.Permissionrequiredforreproductionordisplay.2Mg(s)+O2(g)2MgO(s)2Mg2Mg2++4e-O2+4e-2O2-Oxidationhalf-reaction(losee-)Reductionhalf-reaction(gaine-)19.1Electrochemicalprocessesareoxidation-reductionreactionsinwhich:•theenergyreleasedbyaspontaneousreactionisconvertedtoelectricityor•electricalenergyisusedtocauseanonspontaneousreactiontooccur0 0 2+ 2-OxidationnumberThechargetheatomwouldhaveinamolecule(oranioniccompound)ifelectronswerecompletelytransferred.1. Freeelements(uncombinedstate)haveanoxidationnumberofzero.Na,Be,K,Pb,H2,O2,P4=02. Inmonatomicions,theoxidationnumberisequaltothechargeontheion.Li+,Li=+1;Fe3+,Fe=+3;O2-,O=-23. Theoxidationnumberofoxygenis usually–2.InH2O2andO22-itis–1.4.44. Theoxidationnumberofhydrogenis+1exceptwhenitisbondedtometalsinbinarycompounds.Inthesecases,itsoxidationnumberis–1.6.Thesumoftheoxidationnumbersofalltheatomsinamoleculeorionisequaltothechargeonthemoleculeorion.5. GroupIAmetalsare+1,IIAmetalsare+2andfluorineisalways–1.HCO3-O=-2 H=+13x(-2)+1+?=-1C=+4OxidationnumbersofalltheatomsinHCO3-?4.4BalancingRedoxEquations19.11. Writetheunbalancedequationforthereactionionionicform.TheoxidationofFe2+toFe3+byCr2O72-inacidsolution?Fe2++Cr2O72-Fe3++Cr3+2. Separatetheequationintotwohalf-reactions.Oxidation:Cr2O72-Cr3++6 +3Reduction:Fe2+Fe3++2 +33. BalancetheatomsotherthanOandHineachhalf-reaction.Cr2O72-2Cr3+BalancingRedoxEquations4. Forreactionsinacid,addH2OtobalanceOatomsandH+tobalanceHatoms.Cr2O72-2Cr3++7H2O14H++Cr2O72-2Cr3++7H2O5. Addelectronstoonesideofeachhalf-reactiontobalancethechargesonthehalf-reaction.Fe2+Fe3++1e-6e-+14H++Cr2O72-2Cr3++7H2O6. Ifnecessary,equalizethenumberofelectronsinthetwohalf-reactionsbymultiplyingthehalf-reactionsbyappropriatecoefficients.6Fe2+6Fe3++6e-6e-+14H++Cr2O72-2Cr3++7H2O19.1BalancingRedoxEquations7. Addthetwohalf-reactionstogetherandbalancethefinalequationbyinspection.The number of electrons on both sides must cancel.6e-+14H++Cr2O72-2Cr3++7H2O6Fe2+6Fe3++6e-Oxidation:Reduction:14H++Cr2O72-+6Fe2+6Fe3++2Cr3++7H2O8. Verifythatthenumberofatomsandthechargesarebalanced.14x1–2+6x2=24=6x3+2x319.19. Forreactionsinbasicsolutions,addOH-toboth sidesoftheequationforeveryH+thatappearsinthefinalequation.GalvanicCells19.2spontaneousredoxreactionanodeoxidationcathodereductionGalvanicCells19.2Thedifferenceinelectricalpotentialbetweentheanodeandcathodeiscalled:• cell voltage• electromotive force (emf)• cell potentialCellDiagramZn(s)+Cu2+(aq)Cu(s)+Zn2+(aq)[Cu2+]=1M&[Zn2+]=1MZn(s)|Zn2+(1M)||Cu2+(1M)|Cu(s)anode cathodeStandardReductionPotentials19.3Zn(s)|Zn2+(1M)||H+(1M)|H2(1atm)|Pt(s)2e-+2H+(1M)H2(1atm)Zn(s)Zn2+(1M)+2e-Anode(oxidation):Cathode(reduction):Zn(s)+2H+(1M)Zn2++H2(1atm)StandardReductionPotentials19.3Standard reduction potential (E0)isthevoltageassociatedwithareduction reactionatanelectrodewhenallsolutesare1Mandallgasesareat1atm.E0=0VStandardhydrogenelectrode(SHE)2e-+2H+(1M)H2(1atm)ReductionReaction19.3E0=0.76VcellStandard emf (E0 )cell0.76V=0 - EZn /Zn 02+EZn /Zn=-0.76V02+Zn2+(1M)+2e-ZnE0=-0.76VE0=EH /H - EZn /Zn cell0 0+2+2StandardReductionPotentialsE0=Ecathode - Eanodecell0 0Zn(s)|Zn2+(1M)||H+(1M)|H2(1atm)|Pt(s)StandardReductionPotentials19.3Pt(s)|H2(1atm)|H+(1M)||Cu2+(1M)|Cu(s)2e-+Cu2+(1M)Cu(s)H2(1 atm)2H+(1M)+2e-Anode(oxidation):Cathode(reduction):H2(1atm)+Cu2+(1M)Cu(s)+2H+(1M)E0=Ecathode - Eanodecell0 0E0=0.34VcellEcell=ECu /Cu–EH /H 2+ +20 0 00.34=ECu /Cu-002+ECu