1. 14-1 Electroanalytical chemistry Quantitative methods based on electrical properties when solution is part of an electrochemical cell §Low detection limits §Stoichiometry §Rate of charge transfer §Rate of mass transfer §Absorption §Equilibrium constants of reactions Oxidation state specific Activities rather than concentrations

2. 14-2 Electroanalytical methods Electrochemical cells Potentials in cells Electrode potentials Calculation of cell potentials Types of methods Electrochemical cells §Electrodes in electrolyte solution §Electrodes connected externally §Electrolyte in solution permit ion transfer

3. 14-3 Oxidation and Reduction Primary mechanism for §Batteries §Production of metals from ores Oxidation -reduction occurs simultaneously §For atoms and monatomic ions, loss or gain of electrons §For covalently bonded material can experience bond breaking Used to keep track of electrons in molecule

4. 14-4 Oxidation State Accounts for net charge of molecule Sum of atomic oxidation state comprise molecular state §NaCl: Na + and Cl - §MnO 4 - : Mn 7+ and 4O 2- For free elements each element is assigned an oxidation state of 0 §Hg §Cl 2 §P 4 For monotonic ion, oxidation state is the charge §Cl -, Pu 4+

5. 14-5 Oxidation State Group 1 (IA) elements (Li, Na, K, Rb, Cs, and Fr) are 1+, H can be 1- §ionic hydrides (H with very active metals) àNaH, LiH àLiAlH 4, NaBH 4 Group 2 (IIA) elements (Be, Mg, Ca, Sr, Ba, and Ra) are 2+ Oxygen is usually 2- §Exceptions with oxygen-oxygen bonds àH 2 O 2, Na 2 O 2 : O oxidation state = 1- àKO 2 : O oxidation state = 1/2- àOF 2 : O oxidation state = 2+

6. 14-6 Periodic Variations of Oxidation State constant 1 2 3 Steps of 1 45 6-12 Steps of 2 13-17 18 Mainly 3+

7. 14-7 Oxidizing AgentsReducing Agents F 2 F - Cl 2 Cl - Br 2 Br - Ag + Ag I 2 I - Cu 2+ Cu H + H 2 Fe 2+ Fe Zn 2+ Zn Al 3+ Al Na + Na Oxidizing and Reducing Agents weak Strong

8. 14-8 Redox Reactions Zn + Cu 2+ Zn 2+ + Cu §Zn is oxidized, Cu is reduced §Transfer of electrons from one metal to another May not involved charge species §C + O 2 CO 2 Oxidation agent oxidizes another species and is reduced Reduction agent reduces another species and is oxidized

9. 14-9 Balancing Redox Equations Balancing can be accomplished through examining ion- electron half reactions §H + + NO 3 - + Cu 2 O Cu 2+ + NO + H 2 O Identify reduced and oxidized species §Cu 2 O to Cu 2+ (1+ to 2+): oxidized §NO 3 - to NO (5+ to 2+): reduced Balance oxidized/reduced atoms §Cu 2 O 2Cu 2+ Add electrons to balance redox of element §Cu 2 O 2Cu 2+ + 2e - §NO 3 - + 3e - NO

10. 14-10 Balancing Redox Equations Add H + (or OH - ) to balance charge of reaction §2H + + Cu 2 O 2Cu 2+ + 2e - §4 H + + NO 3 - + 3e - NO Add water to balance O and H, then balance other atoms if needed §2H + + Cu 2 O 2Cu 2+ + 2e - + H 2 O §4 H + + NO 3 - + 3e - NO + 2 H 2 O Multiple equations to normalize electrons §3(2H + + Cu 2 O 2Cu 2+ + 2e - + H 2 O) §2(4 H + + NO 3 - + 3e - NO + 2 H 2 O)

11. 14-11 Balancing Equations Add the reactions together §14H + + 2NO 3 - + 3Cu 2 O 6Cu 2+ +2NO +7 H 2 O Important for reactions involving metal with multiple oxidation states Disproportionation Some elements with intermediate states can react to form species with different oxidation states Species acts as both oxidation and reduction agent §2 Pu 4+ Pu 3+ + Pu 5+

12. 14-12 Electrochemistry Chemical transformations produced by electricity §Corrosion §Refining Electrical Units §Coulomb (C) àCharge on 6.25 x 10 18 electrons §Amperes (A) àElectric current àA=1C/sec

13. 14-13 Electrochemistry Volt (V) §Potential driving current flow §V= 1 J/C Ohm’s law   = IR   = potential, I =current, and R=resistance symbolunitrelationships ChargeqCoulomb (C) CurrentIAmpere (A) I=q/t (t in s) Potential  Volt (V)  =IR PowerPWatt (W) P=  I Energy EJoule (J)Pt=  It=  q ResistanceROhm (  )R=  /I

14. 14-14 Electrolysis Production of a chemical reaction by means of an electric current §2 H 2 O 2H 2 + O 2 Cathode §Electrode at which reduction occurs §Cations migrate to cathode àCu 2+ + 2e - Cu Anode §Electrode at which oxidation occurs §Anions migrate to anode à2Cl - Cl 2 + 2e -

15. 14-15 Electrolysis Redox depends upon tendencies of elements or compounds to gain or lose electrons §electrochemical series àLists of elements or compounds àHalf cell potentials Related to periodic tendencies

16. 14-16 Electrolysis of CuCl 2 C electrode Cu Plating on C electrode C electrode Cl 2 Anode: 2Cl - ->Cl 2 +2e - Cathode: Cu 2+ +2e - ->Cu

17. 14-17 NaCl Solutions Dilute NaCl solution §anode: 2 H 2 O O 2 + 4H + + 4e - §cathode: 2 H 2 O + 2e - H 2 + 2OH - Concentrated NaCl (Brines) §anode: 2Cl - Cl 2 + 2e - §cathode: 2 H 2 O + 2e - H 2 + 2OH - Molten Salt §anode: 2Cl - Cl 2 + 2e - §cathode: Na + + e - Na §Na metal produced by electrolysis of NaCl and Na 2 CO 3 àLower melting point than NaCl

18. 14-18 Faraday Laws In 1834 Faraday demonstrated that the quantities of chemicals which react at electrodes are directly proportional to the quantity of charge passed through the cell 96487 C is the charge on 1 mole of electrons = 1F (faraday)

19. 14-19 Faraday Laws Cu(II) is electrolyzed by a current of 10A for 1 hr between Cu electrode §anode: Cu Cu 2+ + 2e - §cathode: Cu 2+ + 2e - Cu §Number of electrons à(10A)(3600 sec)/(96487 C/mol) = 0.373 F à0.373 mole e - (1 mole Cu/2 mole e - ) = 0.186 mole Cu

20. 14-20 Electrochemical cell

21. 14-21 Conduction in a cell Charge is conducted §Electrodes §Ions in solution §Electrode surfaces àOxidation and reduction àOxidation at anode àReduction at cathode Reaction can be written as half-cell potentials

22. 14-22 Half-cell potentials Standard potential  Defined as  °=0.00V àH 2 (atm) 2 H + (1.000M) + 2e - Cell reaction for §Zn and Fe 3+/2+ at 1.0 M §Write as reduction potentials  Fe 3+ + e - Fe 2+  °=0.77 V  Zn 2+ + 2e - Zn  °=-0.76 V §Fe 3+ is reduced, Zn is oxidized

23. 14-23 Half-Cell Potentials Overall  2Fe 3+ +Zn 2Fe 2+ + Zn 2+  °=0.77+0.76=1.53 V Half cell potential values are not multiplied Application of Gibbs If work is done by a system  ∆G = -  °nF (n= e - ) Find ∆G for Zn/Cu cell at 1.0 M  Cu 2+ + Zn Cu + Zn 2+   °=1.10 V §2 moles of electrons (n=2) à∆G =-2(96487C/mole e - )(1.10V) à∆G = -212 kJ/mol

24. 14-24 Reduction Potentials Electrode Couple"E0, V" Na+ + e- --> Na-2.7144 Mg2+ + 2e- --> Mg-2.3568 Al3+ + 3e- --> Al-1.676 Zn2+ + 2e- --> Zn-0.7621 Fe2+ + 2e- --> Fe-0.4089 Cd2+ + 2e- --> Cd-0.4022 Tl+ + e- --> Tl-0.3358 Sn2+ + 2e- --> Sn-0.141 Pb2+ + 2e- --> Pb-0.1266 2H+ + 2e- --> H2(SHE)0 S4O62- + 2e- --> 2S2O32-0.0238 Sn4+ + 2e- --> Sn2+0.1539 SO42- + 4H+ + 2e- --> H2O + H2SO3(aq)0.1576 Cu2+ + e- --> Cu+0.1607 S + 2H+ + 2e- --> H2S0.1739 AgCl + e- --> Ag + Cl-0.2221 Saturated Calomel (SCE)0.2412 UO22+ + 4H+ + 2e- --> U4+ + 4H2O0.2682

25. 14-25 Reduction Potentials Hg2Cl2 + 2e- --> 2Cl- + 2Hg0.268 Bi3+ + 3e- --> Bi0.286 Cu2+ + 2e- --> Cu0.3394 Fe(CN)63- + e- --> Fe(CN)64-0.3557 Cu+ + e- --> Cu0.518 I2 + 2e- --> 2I-0.5345 I3- + 2e- --> 3I-0.5354 H3AsO4(aq) + 2H+ + 2e- -->H3AsO3(aq) + H2O0.5748 2HgCl2 + 4H+ + 2e- -->Hg2Cl2 + 2Cl-0.6011 Hg2SO4 + 2e- --> 2Hg + SO42-0.6152 I2(aq) + 2e- --> 2I-0.6195 O2 + 2H+ + 2e- --> H2O2(l)0.6237 O2 + 2H+ + 2e- --> H2O2(aq)0.6945 Fe3+ + e- --> Fe2+0.769 Hg22+ + 2e- --> Hg0.7955 Ag+ + e- --> Ag0.7991 Hg2+ + 2e- --> Hg0.8519 2Hg2+ + 2e- --> Hg22+0.9083 NO3- + 3H+ + 2e- -->HNO2(aq) + H2O0.9275

26. 14-26 Reduction Potentials VO2+ + 2H+ + e- --> VO2+ + H2O1.0004 HNO2(aq) + H+ + e- --> NO + H2O1.0362 Br2(l) + 2e- --> 2Br-1.0775 Br2(aq) + 2e- --> 2Br-1.0978 2IO3- + 12H+ + 10e- -->6H2O + I21.2093 O2 + 4H+ + 4e- --> 2H2O1.2288 MnO2 + 4H+ + 2e- -->Mn2+ + 2H2O1.1406 Cl2 + 2e- --> 2Cl-1.3601 MnO4- + 8H+ + 5e- -->4H2O + Mn2+1.5119 2BrO3- + 12H+ + 10e- -->6H2O + Br21.5131

27. 14-27 Nernst Equation Compensated for non unit activity (not 1 M) Relationship between cell potential and activities aA + bB +ne - cC + dD At 298K 2.3RT/F = 0.0592 What is potential of an electrode of Zn(s) and 0.01 M Zn 2+ Zn 2+ +2e - Zn  °= -0.763 V activity of metal is 1

28. 14-28 Electrodes SHE (Standard Hydrogen Electrode) §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 Ag/AgCl electrode §AgCl (s) + e- «Cl- + Ag(s) §Ecell = +0.20 V vs. SHE Calomel electrode §Hg 2 Cl 2 (s) + 2e- «2Cl- + 2Hg(l) §Ecell = +0.24 V vs.SHE

29. 14-29 IR drop Force needed to overcome resistance of ion movement §Follows Ohm’s law §Increase potential required to operate cell §E Cell =E cathode -E anode -IR For a Cd/Cu cell at 4  find potential needed for 0.1 A Cu2+ + 2e- --> Cu0.3394 Cd2+ + 2e- --> Cd-0.4022 Cu 2+ +Cd Cu+Cd 2+: E cell =0.3394-(-0.4022)-4*0.1=0.3416 V

30. 14-30 Polarization E Cell =E cathode -E anode -IR §predicts linear relationship between cell voltage and current §Deviation due to polarity of cell àCan occur at either electrode Due to limitations of reaction at surface of electrode §Mass transfer §Concentration §Reaction intermediates §Physical processes àSorption àCrystallization

31. 14-31 Methods