Notes Ch 23

CH 23Electrochemistry23.1 Electrochemical cellsTypes of electrochemical cellsGalvanic or Voltaic The ‘spontaneous’ reaction.Produces electrical energy.Electrolytic Non-spontaneous reaction.Requires electrical energy to occur.For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.Voltaic cellsThere are two general ways to conduct an oxidation-reduction reactionMixing oxidant and reductant together Cu2+ + Zn(s) Cu(s) + Zn2+ This approach does not allow forcontrol of the reaction.Voltaic cellsElectrochemical cells Each half reaction is put in a separate ‘half cell.’ They can then be connected electrically.This permits better control over the system.Spontaneous ReactionsWill occur if the anode metal is above the cathode metal in the Activity Series chart (pg 678)Voltaic CellAllessandro Volta (1745-1827) invented the first electrochemical cell- this type was called the voltaic cell.It is a spontaneous reaction- he layered Cu and Zn plates, separated by cardboard: Cu plate had reduction occur, Zn plate had oxidation occurVoltaic CellA half-cell is one part of the voltaic cell where either oxidation or reduction occurs. A half cell consists of a metal strip immersed in a solution of it’s ionsVoltaic cellsVoltaic cellsVoltaic cellsSalt bridgeAllows ion migration in solution but prevents extensive mixing of electrolytes.It can be a simple porous disk or a gel saturated with a non-interfering, strong electrolyte like KCl.Voltaic cellsVoltaic cellsVoltaic CellCell diagramsRather than drawing an entire cell, a type of shorthand can be used.For our copper - zinc cell, it would be: Zn | Zn2+ (1M) || Cu2+ (1M) | The anode is always on the left.| = boundaries between phases||= salt bridgeOther conditions like concentration are listed just after each species.Dry CellVoltaic cell where the electrolyte is a paste- not a solutionExample: flashlight battery ( pg 681)Not a true batteryOuter Zn case is anode (oxidation)Carbon (graphite core) rod in center is cathode- but actually reduction occurs w/MnO2 found in pasteSalt bridge is not needed because of paste prevent cell contents from mixingAlkaline batteries use KOH in paste and this makes it last longer and keeps voltage upLead Storage BatteryA battery is a group of cells connected togetherA car battery is 6 cells producing 2V each for a total of 12 VThe cathode is lead(IV) oxide and the anode is Pb. Dilute sulfuric acid is the electrolyteOverall reaction is: Pb(s) + PbO2(s) + 2H2SO4(aq)-----2PbSO4(s) + 2 H2O(l)Now you write the half reactions that occur at each electrode!!Lead BatteryCar battery’s are recharged when the car runs- the reaction occurs in reverse- but this reverse reaction is nonspontaneous and so the car’s generator supplies the energy to drive the reaction.Eventually the battery dies- electrodes lose so much PbSO4 which can fall to the bottom of the batteryFuel CellIdea here is to have a renewable electrode so electrodes don’t wear outA fuel is used for the oxidationSimplest is the Hydrogen-oxygen fuel cell- Oxygen is fiels for cathode (reduction), and hydrogen is fuel for anode (oxidation)Overall reaction: 2H2(g) + O2(g)—2H2O(l) Fuel CellYou write the anode and cathode half-cell reactions.Advantage: cheap fuel, only “pollutant”- water which is drinkableUsed in spacecraft and some military applications- some cars; expensive and takes room.23.2 Half-cells and Cell PotentialsElectrode potentialsA measure of how willing a species is to gain or lose electrons.Standard potentialsPotential of a cell acting as a cathodecompared to a standard hydrogen electrode.Values also require other standard conditions.Standard hydrogen electrodeHydrogen electrode (SHE) The ultimate reference electrode.H2 is constantly bubbled into a 1 M HCl solution Eo = 0.000 000 VAll other standard potentials are then reported relative to SHE.Electrode potentialsStandard potentials are defined using specific concentrations.All soluble species are at 1 MSlightly soluble species must be at saturation.Any gas is constantly introduced at 1 atmAny metal must be in electrical contactOther solids must also be present and in contact.Electrode potentialsThe standard potential for: Cu2+ + 2e- Cu (s)is +0.337V.This means that: If a sample of copper metal is placed in a 1 M Cu2+ solution, we’ll measure a value of 0.337V if compared to: 2H+ + 2e- H2 (g) (1 M) (1atm)Half reactionsA common approach for listing species that undergo REDOX is as half-reactions.For 2Fe3+ + Zno(s) = 2Fe2+ + Zn2+ Fe3+ + e- Fe2+(reduction) Zno(s) Zn2+ + 2e- (oxidation)You’ll find this approach useful for a number of reasons.Half reactionsTables are available which list half reactions as either oxidations or reductions.( Pg 688)Will provide Standard Eo values to help predict reactions and equilibria.Other species that participate in the reaction.Show the relative ability to gain or loss electrons.Cell potentialsOne thing that we would like to know is the spontaneous direction for a reaction.This requires that we determine the Ecell.Since our standard potentials (E o) are commonly listed as reductions, we’ll base our definitions on that.Ecell = Ehalf-cell of reduction - Ehalf-cell of oxidation Eocell = Eohalf-cell of reduction - Eohalf-cell of oxidationCell potentialsYou know that both an oxidation and a reduction must occur. One of your half reactions must be reversed.The spontaneous or galvanic direction for a reaction is the one where Ecell is a positive value.The half reaction with the largest E value will proceed as a reduction.The other will be reversed - oxidation. Cell potentialsFor our copper - zinc cell at standard conditions: Cu2+ + 2e- Cu (s)+0.34 V Zn2+ + 2e- Zn (s) -0.763 VEcell 1.03 VSpontaneous reaction at standard conditions: Cu2+ + Zn (s) Cu (s) + Zn2+Concentration dependency of E Eo values are based on standard conditions.The E value will vary if any of the concentrations vary from standard conditions.This effect can be experimentally determined by measuring E versus a standard (indicator) electrode.Theoretically, the electrode potential can be determined by the Nernst equation.Calculation of cell potentialsTo determine the Voltaic Ecell at standard conditions using reduction potentials:Ecell = E ohalf-cell of reduction - E ohalf-cell of oxidation Where Ehalf-cell of reduction - half reaction with the larger , or least negative E o value.Ehalf-cell of oxidation - half reaction with the smaller or more negative E o value.Calculation of cell potentialsSteps in determining the spontaneous direction and E of a cell at standard conditions Calculate the E for each half reaction.The half reaction with the largest or least negative E value will proceed as a reduction.Calculate EcellCalculation of cell potentialsExampleDetermine the spontaneous direction and Ecell for the following system. Pb | Pb2+ (1M) || Sn2+ (1M) | SnHalf reaction Eo Pb2+ + 2e- Pb-0.13 V Sn2+ + 2e- Sn-0.14 VNote: The above cell notation may or may not be correct.Calculation of cell potentials anode: Pb2+ + 2e- Pb -0.13 Vcathode:Sn2+ + 2e- Sn -0.14 VEcell = E ohalf-cell of reduction - E ohalf-cell of oxidation (-0.14) – (-0.13)= -0.1 V so not spontaneousCalculation of cell potentialsCalculate the E0cell to determine if the following cell is spontaneous as writtenNi(s) + Fe2+(aq) -----Ni2+(aq) + Fe(s) Calculation of cell potentialsWrite each half reaction and look up the reduction cell potentialOxidation: Ni----Ni2+ + 2e--0.25 VReduction:Fe2+ +2e-------Fe -0.44 VE0cell= (-0.44)-(-0.25)= -0.19VBecause the potential is negative, it is not spontaneousCalculation of cell potentialsA voltaic cell is constructed using the following half-reactions:Ag+ + 1e------Ag +0.80Cu2++2e------Cu +0.34Determine the cell reaction and the standard cell potentialCalculation of cell potentialsFirst determine which will undergo reduction and which will oxidizeAg+ + 1e------Ag +0.80 reduction (larger)Cu2++2e------Cu +0.34 oxide (smaller)E0cell= +0.80 -(+0.34)= +0.46 V 2Ag+ + Cu ------2Ag + Cu2+ 23.3 Electrolytic cellsElectrolytic CellsAn electrochemical cell that uses electrical energy to cause a chemical reaction ( reverse of a voltaic cell)Uses energy to drive a non-spontaneous redox reactionComparing Voltaic and Electrolytic CellsComparing Electrolytic & Voltaic CellsIn both, electrons flow from anode to cathodeIn both, reduction occurs at the cathode; oxidation occurs at the anodeIn voltaic cell- reaction is spontaneous; in electrolytic it is non-spontaneous- the electrons are pushed by an outside source ( electricity)Comparing Electrolytic & Voltaic CellsIn an electrolytic cell the anode is +, in a voltaic remember it is –In an electrolytic cell the cathode is -, in a voltaic is is +Electrolysis of WaterSplitting of water using electricity- this is done by an electrolytic cellNeed an electrolyte in water- complete the circuit- usually it is sulfuric acid or sodium hydroxideThe products are hydrogen gas ( at cathode) and oxygen ( at anode)Electrolysis of WaterAt Cathode ( - electrode) Reduction: 2H2O(l) +2e- -----H2(g) + 2OH-(aq)Notice- the solution around the cathode will become basicAt Anode (+ electrode) Oxidation: 2H2O(l) -----O2(g) + 4H+(aq) +4e-Notice- the solution around the anode will become acidicElectrolysis of WaterOverall net reaction:2H2O(l) electrolysis 2H2(g) + O2(g)Extra H+ and OH- ions combine to reform water and so are not included I nt ehfinal equationElectrolysis of BrineBrine is a concentrated aqueous solution of saltStart with NaCl and H2OAt Cathode (- electrode) reduction: 2H2O(l) ------H2(g)+2OH-(aq)Notice: H gets reduced, not Na because water is more easily reduced- the Na becomes a spectator ion which will combine with another spectator ion OH- Electrolysis of BrineAt Anode (+ electrode) Oxidation: 2Cl-(aq) ------Cl2(g)+2e-Overall Reaction:2NaCl(aq) + 2H2O(l) -----Cl2(g) + H2 (g)+2NaOH(aq) The chlorine is collected and used as disinfectantsHydrogen can be collected, and NaOH can also be recovered and usedElectrolysis of Molten Sodium ChlorideSodium is used in vapor lamps and as a coolant in some nuclear reactorsChlorine is used as to produce pesticides, manufacture polyvinyl chloride and as a disinfectantAt Cathode (- electrode) Reduction: 2Na+(l) + 2 e- ------2Na(l)Electrolysis of Molten Sodium ChlorideAt Anode (+ electrode) 2Cl-(l) ------Cl2(g) + 2e-Overall reaction:2NaCl(l) electrolysis 2Na(l) + Cl2(g) ElectroplatingThis is the deposition of a thin, very thin ( 10-5cm) layer of metal on an object in an electrolytic cellGold, silver, copper, nickel and chromium are the most commonly used to plate on other objectsElectroplatingExample: to plate silver on an objectMake Ag the anode and the object the cathodeAt the anode the silver becomes oxidized ( Ag+) and these ions migrate to the cathode where they become reduced on the object ( Ag+ + e- -----Ag)ElectroplatingMakes things resist corrosion, shiny, more appealing also can be used to make metal molds ( DVD’s or phonographs)In electrowinning impure metals can be purified in electrolytic cellsElectropolishing is another process which removes a thin layer of metal to give a high polish