1. Putting redox reactions to work

2.  Electrons are transferred  Lose Electrons Oxidation  Gain Electrons Reduction

3.  Made of two half-cells  Based upon two half-reactions  Electrons travel between the two half-cells

4.  Also called voltaic cells  Convert chemical energy into electrical energy  Spontaneous

5.  Convert electrical energy into chemical energy  Non-spontaneous

6.  Write the reaction for solid magnesium placed in a copper (II) sulfate solution. Mg (s) + CuSO 4 (aq)  MgSO 4 (aq) + Cu (s)

7.  Balance the reaction using the half-reaction method Mg (s) + CuSO 4 (aq)  MgSO 4 (aq) + Cu (s) 0+2 +6 -2 +2 +6 -2 0

8.  Balance the reaction using the half-reaction method Mg (s) + CuSO 4 (aq)  MgSO 4 (aq) + Cu (s) 0+2 +6 -2 +2 +6 -2 0 Mg (s)  Mg 2+ (aq) + 2e - Cu 2+ (aq) + 2e -  Cu (s)

9.  Balance the reaction using the half-reaction method Mg (s) + CuSO 4 (aq)  MgSO 4 (aq) + Cu (s) 0+2 +6 -2 +2 +6 -2 0 1(Mg (s)  Mg 2+ (aq) + 2e - ) 1(Cu 2+ (aq) + 2e -  Cu (s)) Mg (s)+CuSO 4 (aq)+2e -  MgSO 4 (aq)+Cu (s)+2e -

11.  Anode  Cathode  Salt Bridge  Flow of electrons

12.  Potential (either half-cell or cell) ◦ Pull on the electrons ◦ Electromotive force (emf) ◦ Volt (V) ◦ Joule/Coulomb (J/C)  Voltmeter ◦ Analog ◦ Digital  Potentiometer  Positive potential…spontaneous  Negative potential…nonspontaneous

13.  Standard Reduction Potentials Chart ◦ Only reduction reactions ◦ Must look up the reverse of the oxidation and flip the sign of the potential  Add standard half-cell potentials to get standard cell potential

14. 1(Mg (s)  Mg 2+ (aq) + 2e - ) E˚ ox = +2.37 V 1(Cu 2+ (aq) + 2e -  Cu (s)) E˚ red = +0.342 V Mg (s)+CuSO 4 (aq)+2e -  MgSO 4 (aq)+Cu (s)+2e - E˚ cell = +2.71 V

15.  Oxidation||Reduction  X(s)|X + (aq)||Y + (aq)|Y(s)  Mg(s)|Mg 2+ (aq)||Cu 2+ (aq)|Cu(s) 1(Mg (s)  Mg 2+ (aq) + 2e - ) 1(Cu 2+ (aq) + 2e -  Cu (s)) Mg (s)+CuSO 4 (aq)+2e -  MgSO 4 (aq)+Cu (s)+2e -

16.  Cu(s)|Cu 2+ (aq)||Ag 1+ (aq)|Ag(s) 1(Cu (s)  Cu 2+ (aq) + 2e - ) E˚ ox = -0.342V 2(Ag 1+ (aq) + 1e -  Ag (s)) E˚ red = +0.800V Cu(s)+2Ag 1+ (aq)+2e -  Cu 2+ (aq)+2Ag (s)+2e - E˚ cell = +0.458V

17.  Series of electrochemical cells connected to each other  Completes the circuit  Dry cell ◦ Flashlight battery ◦ Watch battery  Wet Cell ◦ Car battery

18.  Carbon-Zinc Battery ◦ Zinc casing…anode ◦ Carbon rod…cathode ◦ MnO 2 is actually reduced ◦ Alkaline battery…has KOH rather than NH 4 Cl

19.  Carbon-Zinc Battery ◦ Zn(s)  Zn 2+ (aq) + 2e - ◦ 2NH 4 1+ (aq) + 2MnO 2 (s) + 2e -  Mn 2 O 3 (s) + 2NH 3 (g) + H 2 O(l)

20.  Lead-Acid Storage Battery  Pb(s) + PbO 2 (s) + H 2 SO 4 (aq)  PbSO 4 (s) + H2O(l)  Spontaneous & nonspontaneous

21.  Not 1M  Require additional calculations  Can manipulate potential to a particular V

22.  Nernst Equation ◦ E cell = E˚ cell – {(0.0592/n)(logQ)}  n  Q

23.  A 0.500M solution of copper (II) sulfate is reacted with magnesium metal. A 0.750M solution of magnesium sulfate is one of the products. What is the cell potential? ◦ Write two half reactions ◦ Write balanced equation ◦ Determine n ◦ Determine E˚ cell ◦ Use Nernst to solve for E cell  Mg(s)|Mg 2+ (aq)||Cu 2+ (aq)|Cu(s)

24. 1(Mg (s)  Mg 2+ (aq) + 2e - ) E˚ ox = +2.37 V 1(Cu 2+ (aq) + 2e -  Cu (s)) E˚ red = +0.342 V Mg (s)+CuSO 4 (aq)+2e -  MgSO 4 (aq)+Cu (s)+2e - E˚ cell = +2.71 V

25. E cell = 2.71 V – {(0.0592/2)(log([0.75]/[0.5]))} E cell = 2.71 V – {(0.0296)(0.176)} E cell = 2.71 V – 0.00521 E cell = 2.70 V

26.  Cu(s)|Cu 2+ (0.0100M)||Ag 1+ (0.0250M)|Ag(s)