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Fig. 22-1a (p.629) A galvanic electrochemical cell at open circuit.

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Presentation on theme: "Fig. 22-1a (p.629) A galvanic electrochemical cell at open circuit."— Presentation transcript:

1 Fig. 22-1a (p.629) A galvanic electrochemical cell at open circuit

2 1.1 components Conducting electrodes (metal, carbon) External wires (electrons carry current) Ion electrolyte solution (ions carry current) Interfaces or junctions Complete electrical circuit conduction: electrons moving from one electrode to another thru. External wire within solutions, migration of ions carry current at electrode surface, oxidation or reduction reactions occur coupling the electron conduction of electrodes w/ ion conduction of solutions

3 1.2 Galvanic Cells – cell develops spontaneous potential difference Overall: 2Ag + (aq) + Cu(s)  2Ag(s) + Cu 2+ (aq) Half reactions: Cu(s)  Cu 2+ (aq) + 2e - Oxidation Ag + + e -  Ag (s) Reduction Convention: Cathode where reduction occurs Anode where oxidation occurs Galvanic cell: Cu anode (negative) Ag cathode (positive)  Cu2+ = 0.02 M,  Ag+ = 0.02 M, E = V  0 V reaction  equilibrium Potential difference (voltage) is measure of tendency to move to equilibrium

4 Fig. 22-1b (p.629) A galvanic cell doing work

5 1.3 electrolytic cells – require potential difference greater than galvanic potential difference ( to drive away from equilibrium) Overall: 2Ag(s) + Cu 2+ (aq)  2Ag + (aq) + Cu(s) [chemically reversible cell] Half reactions: Ag(s)  Ag + (aq) + 2e - Oxidation Cu 2+ (aq) + 2e -  Cu(s)Reduction electrolytic cell: Ag anode (negative) Cu cathode (positive)

6 Fig. 22-1c (p.629) An electrolytic cell

7 1.4Cell w/o liquid junctions AgCl(s)  Ag+ (aq) + Cl- (aq) H2 (g)  H2 (aq) Cathode: AgCl(s)  Ag + + Cl - Ag + + e -  Ag (s) - Anode: H 2 (aq)  2H + (aq) + 2e - Overall: 2AgCl (s) + H2(g)  2Ag(s) + 2H + + 2Cl - direct reaction of AgCl + H 2 is very slow

8 Fig (p.631) A galvanic cell without a liquid junction

9 1.5 Schematic representation of cells Short-hand cell notation Cu|CuSO 4 (  Cu2+ = ) ||AgNO 3 (  Ag+ = )| Ag | liquid-electrode interface || two phase boundaries, one at each end of the salt bridge convention: anode on left Galvanic cell as written Electrolytic cell if reversed Pt, H 2 (p=1atm)|H+ (0.01M), Cl - (0.01M), AgCl (sat’d) | Ag

10 2.1 Cell potential (difference between anode and cathode potential) E cell = E cathode – E anode when half-reactions written as reduction (electron on left) Example: 2AgCl(s) + H 2 (g)  2Ag (s) + 2H + + 2Cl - Cathode: 2AgCl(s) + 2e -  2Ag (s) + 2Cl - Anode: 2H + + 2e -  H 2 (g) Galvanic cell E cell = E cathode – E anode = E AgCl/Ag – E H+/H2 = V Can’t measure potential on each electrode independently – only differences

11 2.2 Standard reference electrode Standard hydrogen electrode (SHE) Pt, H 2 (p=1.00atm) | H + (  H + = 1.00M)||… - SHE assigned 0.000V - can be cathode or anode {depending on the half cell which it couples with} - Pt does not take part in reaction, coated with a finely divided layer of platinum to provide large surface area - controlled activity of reactants - rarely used for routine measurement

12 Fig (p.637) measurement of the electrode potential for M electrode

13 Alternative references electrodes {which are simple to prepare} a. Calomel electrode Hg|Hg 2 Cl 2 (sat’d), KCl(xM)|| Hg 2 Cl 2 (s) + 2e -  2Cl - + 2Hg (l) E reference depends on  Cl-1 E reference = +0.24V vs. SHE for saturated calomel electrode (SCE)

14 b. Silver-silver chloride electrode Ag|AgCl(sat’d), KCl(xM)|| AgCl(s) + e -  Cl - + Ag (s) - Similar construction to calomel electrode Ag wire coated with AgCl Solution of KCl sat’d with AgCl - Again, E reference depends on  Cl -, and = V vs. SHE - Can be used for uncontrolled temperature (lower temperature coefficient, see Table 23-1) Can be used for temp > 60  C - But Ag reacts with more ions (e.g, proteins), while Hg reacts with few sample components,

15 Fig (p.661) Typical commerical reference electrode a) A saturated calomel electrode, and b) a silver-silver chloride electrode

16 2.3 electrode and standard electrode potential (E and E 0 ) - Definition: potential of electrodes vs. SHE - Electrode potential varies with activity of ion Activity  x =  x [X]  x : activity coefficient, varies with presence of other ions (ionic strength) (see Appendix 2) Note: activity of pure liquid or solid in excess = 1.00 Use pressure (atm) for gases If  = 1.00M, the electrode potential E, becomes standard electrode potential E 0 Appendix 3 Cu e -  Cu(s) E 0 = V 2H e -  H 2 (g) E 0 = V Cd e -  Cd(s) E 0 = V Zn e -  Zn(s) E 0 = V

17 2.4 Nernst equation In general, E and E cell can be calculated for any activity using Nernst equation: pP + qQ + ne -  rR + sS E = E 0 when log quotient is unity E 0 is relative to SHE E 0 is measure of driving force for half-cell reduction

18 3. Electrical double layer* electrons transferred at electrode surface by redox reactions, occurring at solution/electrode interface Electrical double layer formed 1. tightly bound inner layer 2. loosely bound outer layer

19 Fig (p.632) Electrical double layer formed at electrode surface as a result of an applied potentials


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