Introduction to Electrochemistry (Chapter 22) Many different electroanalytical methods: ·fast ·inexpensive ·in situ ·information about oxidation states stoichiometry rates charge transfer equilibrium constants CEM 333 page 10.1
Introduction to Electrochemistry (Chapter 22) Many different electroanalytical methods: • fast • inexpensive • in situ • information about oxidation states stoichiometry rates charge transfer equilibrium constants CEM 333 page 10.1
Electrochemical Cells: Oxidation and reduction (redox)reactions Separate species to prevent direct reaction (Fig 22-1) Voltmeter 1.100V Salt bridge Saturated KCl solutior Zn electrode Cu electrode 0.0100M 0.0100M ZnSO CuS04 solution solution Zn(s)=Zn2+(ag)+2e Cu2+(aq)+2e-=Cu(s) azm2+=0.010 acu2+=0.010 Anode Cathode Most contain· external wires (electrons carry current) ion solutions (ions carry current) interfaces or junctions All contain.complete electrical circuit conducting electrodes (metal,carbon) CEM 333 page 10.2
Electrochemical Cells: Oxidation and reduction (redox) reactions Separate species to prevent direct reaction (Fig 22-1) Most contain • external wires (electrons carry current) • ion solutions (ions carry current) • interfaces or junctions All contain • complete electrical circuit • conducting electrodes (metal, carbon) CEM 333 page 10.2
Electrons transferred at electrode surface at liquid/solid interface Potential difference (voltage)is measure of tendency to move to equilibrium Galvanic cell-cell develops spontaneous potential difference Overall: Zn(s)+Cu2+(aq)>Zn2+(aq)+Cu(s) Zn(s)→Zn2++2e Oxidation Half reactions: Cu2++2e→Cu(S) Reduction Convention: Reduction at Cathode Oxidation at Anode Galvanic cell-Zn anode (negative),Cu cathode (positive) CEM 333 page 10.3
Electrons transferred at electrode surface at liquid/solid interface Potential difference (voltage) is measure of tendency to move to equilibrium Galvanic cell - cell develops spontaneous potential difference Overall: Zn(s) + Cu2+ (aq) ® Zn2+ (aq) + Cu(s) Half reactions: Zn(s) ® Zn2+ + 2e- Oxidation Cu2+ + 2e- ® Cu(s) Reduction Convention: Reduction at Cathode Oxidation at Anode Galvanic cell - Zn anode (negative), Cu cathode (positive) CEM 333 page 10.3
Electrolytic cells-require potential difference greater than galvanic potential difference(to drive away from equilibrium) Zn(s)→Zn2++2e Oxidation Galvanic cell Cu2++2e→Cu(s) Reduction Zn2++2e→Zn(s) Reduction Electrolytic cell Cu(s)>Cu2++2e-Oxidation Electrolytic cell-Zn cathode(positive),Cu anode(negative) Many chemically reversible cells Short-Hand Cell notation: Convention: Anode on Left Zn|ZnSO(0.01 M)CuSO(0.01 M)Cu liquid-liquid interface Galvanic cell as written Electrolytic cell if reversed CEM 333 page 10.4
Electrolytic cells - require potential difference greater than galvanic potential difference (to drive away from equilibrium) Zn(s) ® Zn2+ + 2e- Oxidation Cu2+ + 2e- ® Cu(s) Reduction Galvanic cell Zn2+ + 2e- ® Zn(s) Reduction Cu(s) ® Cu2+ + 2e- Oxidation Electrolytic cell Electrolytic cell - Zn cathode (positive), Cu anode (negative) Many chemically reversible cells Short-Hand Cell notation: Convention: Anode on Left Zn|ZnSO4 (0.01 M)||CuSO4 (0.01 M)|Cu liquid-liquid interface Galvanic cell as written Electrolytic cell if reversed CEM 333 page 10.4
Not all cells have liquid-liquid junctions(Fig 22-3) H2 10 P=1.00atm) 0.01MHC1 saturated with AgCl Ptelectrode(anode Hz(aq)-H+(aq)+e Solid AgCl AgCl(s)->Ag(aq)+CI(aq) H2(g)→H2(aq) Cathode:Ag(aq)+e->Ag(s) Anode: H2(aq)→2H(aq)+2e Overall:2AgCl(s)+H(g)>2Ag(s)+2H++2CI- Pt,H2(p=1atm)H(0.01 M),CI(0.01 M),AgCI(sat'd)Ag CEM 333 page 10.5
Not all cells have liquid-liquid junctions (Fig 22-3) AgCl(s) ® Ag+ (aq) + Cl- (aq) H2 (g) ® H2 (aq) Cathode: Ag + (aq) + e - ® Ag(s) Anode: H2 (aq) ® 2H+ (aq) + 2e - Overall: 2AgCl(s) + H2 (g) ® 2Ag(s) + 2H+ + 2ClPt,H 2 (p = 1atm)|H + (0.01 M),Cl - (0.01 M),AgCl (sat'd)|Ag CEM 333 page 10.5
Electrode Potentials: Cell potential is difference between anode and cathode potential Ecell =Ecathode-Eanode when half-reactions written as reductions Example: 2AgCl(s)+H2(g)>2Ag(s)+2H++2CI- 2AgCI(s)+2e>2Ag(s)+2CI- 2Ht+2e←→H2(g) electrons on left Galvanic cell Ecell=Ecathode-Eanode=+0.46 V Can't measure potential on each electrode independently -only differences CEM 333 page 10.6
Electrode Potentials: • Cell potential is difference between anode and cathode potential Ecell = Ecathode - Eanode when half-reactions written as reductions Example: 2AgCl(s) + H2 (g) ® 2Ag(s) + 2H+ + 2Cl- 2AgCl(s) + 2e- « 2Ag(s) + 2Cl- 2H+ + 2e- « H2 (g) electrons on left Galvanic cell Ecell=Ecathode-Eanode=+0.46 V Can't measure potential on each electrode independently - only differences CEM 333 page 10.6
Standard reference electrode is usually standard hydrogen electrode (SHE) Pt,H2(p=1.00 atm)H(a=1.00 M)Il. Fig 22-5 Voltmeter e- 0.337V Salt bridge H2 gas 0 PH2=1.00 atm aH+=1.00 aw2+=1.00 CEM 333 page 10.7
Standard reference electrode is usually standard hydrogen electrode (SHE) Pt,H2 (p =1.00 atm)|H+ (aH + = 1.00 M)||. Fig 22-5 CEM 333 page 10.7
SHE: assigned 0.000 V can be anode or cathode Pt does not take part in reaction Pt electrode coated with fine particles(Pt black)to provide large surface area cumbersome to operate Alternative reference electrodes: ·Ag/AgCl electrode AgCI(s)+e←CI-+Ag(s) Ecell =+0.20 V vs.SHE ·Calomel electrode Hg2Cl,(s)+2e←→2C1-+2Hg) Ecell =+0.24 V vs.SHE CEM 333 page 10.8
SHE: • assigned 0.000 V • can be anode or cathode • Pt does not take part in reaction • Pt electrode coated with fine particles (Pt black) to provide large surface area • cumbersome to operate Alternative reference electrodes: • Ag/AgCl electrode AgCl(s) + e - « Cl- + Ag(s) Ecell = +0.20 V vs. SHE • Calomel electrode Hg2Cl2 (s) + 2e- « 2Cl - + 2Hg(l) Ecell = +0.24 V vs.SHE CEM 333 page 10.8
Electrode and Standard Electrode Potentials(E and E0): How do we know which way reaction will go spontaneously? Use electrode potentials,E(potential of electrode versus SHE)to find Eanode and Ecathode.Then find Ecell. But electrode potential varies with activity of ion(appendix 2) activity activity coefficient ax=Yx[X] concentration Yx varies with presence of other ions (ionic strength) u=Xzx×2+Yz2+ concentration charge Note:activity of pure liquid or solid in excess=1.00 Note:use pressure (atm)for gases If a=1.00 M.the electrode potential,E,becomes standard electrode potential,E0 CEM 333 page 10.9
Electrode and Standard Electrode Potentials (E and E0): How do we know which way reaction will go spontaneously? Use electrode potentials, E (potential of electrode versus SHE) to find Eanode and Ecathode. Then find Ecell. But electrode potential varies with activity of ion (appendix 2) activity activity coefficient aX = g X ×[X] concentration gX varies with presence of other ions (ionic strength) m = 1 2 [X]Z X 2 +[Y]ZY 2 ( +.) concentration charge Note: activity of pure liquid or solid in excess=1.00 Note: use pressure (atm) for gases If a=1.00 M, the electrode potential, E, becomes standard electrode potential, E0 CEM 333 page 10.9
Appendix 3: Cu2++2e←Cu(S) E0=+0.337V 2H++2e←→H2(g）) E0=+0.000V Cd2++2e->Cd(s) E0=-0.403V Zn2++2e>Zn(s) E0=-0.763V Cell containing Cu/Cu2+and Cd/Cd2+ called couple (1)Cu2++2e->Cu spontaneously forward Cd2++2e->Cd spontaneously backward(Cd->Cd2++2e-) (2)e-flow towards Cu electrode (cathode/positive electrode) e-flow away from Cd electrode(anode/negative electrode) (3)Cu2+good electron acceptor (oxidizing agent) Cd good electron donor (reducing agent) The most positive E or E0 spontaneously forward forming cathode CEM 333 page 10.10
Appendix 3: Cu2+ + 2e- « Cu(s) E 0 = +0.337 V 2H+ + 2e- « H2 (g) E 0 = +0.000 V Cd2+ + 2e- « Cd(s) E 0 = -0.403 V Zn2+ + 2e- « Zn(s) E 0 = -0.763 V Cell containing Cu/Cu2+ and Cd/Cd2+ called couple (1) Cu2++2e-®Cu spontaneously forward Cd2++2e-®Cd spontaneously backward (Cd®Cd2++2e-) (2) e- flow towards Cu electrode (cathode/positive electrode) e- flow away from Cd electrode (anode/negative electrode) (3) Cu2+ good electron acceptor (oxidizing agent) Cd good electron donor (reducing agent) The most positive E or E0 spontaneously forward forming cathode CEM 333 page 10.10