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Electroanalytical chemistry
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  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. 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. 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. Oxidation State • Accounts for net charge of molecule • Sum of atomic oxidation state comprise molecular state • NaCl: Na+ and Cl- • MnO4-: Mn7+ and 4O2- • For free elements each element is assigned an oxidation state of 0 • Hg • Cl2 • P4 • For monotonic ion, oxidation state is the charge • Cl-, Pu4+

  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) • NaH, LiH • LiAlH4, NaBH4 • Group 2 (IIA) elements (Be, Mg, Ca, Sr, Ba, and Ra) are 2+ • Oxygen is usually 2- • Exceptions with oxygen-oxygen bonds • H2O2, Na2O2: O oxidation state = 1- • KO2: O oxidation state = 1/2- • OF2: O oxidation state = 2+

  6. Periodic Variations of Oxidation State 1 18 13-17 2 6-12 3 4 5 Steps of 1 constant Steps of 2 Mainly 3+

  7. Oxidizing and Reducing Agents Oxidizing Agents Reducing Agents F2F- Cl2Cl- Br2Br- Ag+ Ag I2I- Cu2+ Cu H+ H2 Fe2+ Fe Zn2+Zn Al3+Al Na+Na Strong weak

  8. Redox Reactions • Zn + Cu2+ <--> Zn2+ + Cu • Zn is oxidized, Cu is reduced • Transfer of electrons from one metal to another • May not involved charge species • C + O2 <--> CO2 • Oxidation agent oxidizes another species and is reduced • Reduction agent reduces another species and is oxidized

  9. Balancing Redox Equations • Balancing can be accomplished through examining ion-electron half reactions • H+ + NO3- + Cu2O <--> Cu2+ + NO + H2O • Identify reduced and oxidized species • Cu2O to Cu2+ (1+ to 2+): oxidized • NO3- to NO (5+ to 2+): reduced • Balance oxidized/reduced atoms • Cu2O <--> 2Cu2+ • Add electrons to balance redox of element • Cu2O <--> 2Cu2+ + 2e- • NO3- + 3e- <--> NO

  10. Balancing Redox Equations • Add H+ (or OH-) to balance charge of reaction • 2H+ + Cu2O <--> 2Cu2+ + 2e- • 4 H+ + NO3- + 3e- <--> NO • Add water to balance O and H, then balance other atoms if needed • 2H+ + Cu2O <--> 2Cu2+ + 2e- + H2O • 4 H+ + NO3- + 3e- <--> NO + 2 H2O • Multiple equations to normalize electrons • 3(2H+ + Cu2O <--> 2Cu2+ + 2e- + H2O) • 2(4 H+ + NO3- + 3e- <--> NO + 2 H2O)

  11. Balancing Equations • Add the reactions together • 14H+ + 2NO3- + 3Cu2O <--> 6Cu2++2NO +7 H2O • 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 Pu4+ <--> Pu3+ + Pu5+

  12. Electrochemistry • Chemical transformations produced by electricity • Corrosion • Refining • Electrical Units • Coulomb (C) • Charge on 6.25 x 1018 electrons • Amperes (A) • Electric current • A=1C/sec

  13. Electrochemistry • Volt (V) • Potential driving current flow • V= 1 J/C • Ohm’s law •  = IR •  = potential, I =current, and R=resistance symbol unit relationships Charge q Coulomb (C) Current I Ampere (A) I=q/t (t in s) Potential e Volt (V) e=IR Power P Watt (W) P= eI Energy E Joule (J) Pt= eIt= eq Resistance R Ohm (W) R= e/I

  14. Electrolysis • Production of a chemical reaction by means of an electric current • 2 H2O <--> 2H2 + O2 • Cathode • Electrode at which reduction occurs • Cations migrate to cathode • Cu2+ + 2e- <--> Cu • Anode • Electrode at which oxidation occurs • Anions migrate to anode • 2Cl- <-->Cl2 + 2e-

  15. Electrolysis • Redox depends upon tendencies of elements or compounds to gain or lose electrons • electrochemical series • Lists of elements or compounds • Half cell potentials • Related to periodic tendencies

  16. Electrolysis of CuCl2 Anode: 2Cl-->Cl2+2e- Cathode: Cu2++2e-->Cu C electrode C electrode Cu Plating on C electrode Cl2

  17. NaCl Solutions • Dilute NaCl solution • anode: 2 H2O <--> O2 + 4H+ + 4e- • cathode: 2 H2O + 2e- <-->H2 + 2OH- • Concentrated NaCl (Brines) • anode: 2Cl- <--> Cl2 + 2e- • cathode: 2 H2O + 2e- <-->H2 + 2OH- • Molten Salt • anode: 2Cl- <--> Cl2 + 2e- • cathode: Na+ + e- <-->Na • Na metal produced by electrolysis of NaCl and Na2CO3 • Lower melting point than NaCl

  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. Faraday Laws • Cu(II) is electrolyzed by a current of 10A for 1 hr between Cu electrode • anode: Cu <--> Cu2+ + 2e- • cathode: Cu2+ + 2e- <--> Cu • Number of electrons • (10A)(3600 sec)/(96487 C/mol) = 0.373 F • 0.373 mole e- (1 mole Cu/2 mole e-) = 0.186 mole Cu

  20. Electrochemical cell

  21. Conduction in a cell • Charge is conducted • Electrodes • Ions in solution • Electrode surfaces • Oxidation and reduction • Oxidation at anode • Reduction at cathode • Reaction can be written as half-cell potentials

  22. Half-cell potentials • Standard potential • Defined as °=0.00V • H2(atm) <--> 2 H+ (1.000M) + 2e- • Cell reaction for • Zn and Fe3+/2+ at 1.0 M • Write as reduction potentials • Fe3+ + e- <--> Fe2+ °=0.77 V • Zn2+ + 2e- <-->Zn °=-0.76 V • Fe3+ is reduced, Zn is oxidized

  23. Half-Cell Potentials • Overall • 2Fe3+ +Zn <--> 2Fe2+ + Zn2+ °=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 • Cu2+ + Zn <--> Cu + Zn2+ °=1.10 V • 2 moles of electrons (n=2) • ∆G =-2(96487C/mole e-)(1.10V) • ∆G = -212 kJ/mol

  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- --> H2S 0.1739 AgCl + e- --> Ag + Cl- 0.2221 Saturated Calomel (SCE) 0.2412 UO22+ + 4H+ + 2e- --> U4+ + 4H2O 0.2682

  25. Reduction Potentials Hg2Cl2 + 2e- --> 2Cl- + 2Hg 0.268 Bi3+ + 3e- --> Bi 0.286 Cu2+ + 2e- --> Cu 0.3394 Fe(CN)63- + e- --> Fe(CN)64- 0.3557 Cu+ + e- --> Cu 0.518 I2 + 2e- --> 2I- 0.5345 I3- + 2e- --> 3I- 0.5354 H3AsO4(aq) + 2H+ + 2e- -->H3AsO3(aq) + H2O 0.5748 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- --> Hg 0.7955 Ag+ + e- --> Ag 0.7991 Hg2+ + 2e- --> Hg 0.8519 2Hg2+ + 2e- --> Hg22+ 0.9083 NO3- + 3H+ + 2e- -->HNO2(aq) + H2O 0.9275

  26. Reduction Potentials VO2+ + 2H+ + e- --> VO2+ + H2O 1.0004 HNO2(aq) + H+ + e- --> NO + H2O 1.0362 Br2(l) + 2e- --> 2Br- 1.0775 Br2(aq) + 2e- --> 2Br- 1.0978 2IO3- + 12H+ + 10e- -->6H2O + I2 1.2093 O2 + 4H+ + 4e- --> 2H2O 1.2288 MnO2 + 4H+ + 2e- -->Mn2+ + 2H2O 1.1406 Cl2 + 2e- --> 2Cl- 1.3601 MnO4- + 8H+ + 5e- -->4H2O + Mn2+ 1.5119 2BrO3- + 12H+ + 10e- -->6H2O + Br2 1.5131

  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 Zn2+ • Zn2+ +2e- <--> Zn °= -0.763 V • activity of metal is 1

  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 • Hg2Cl2 (s) + 2e- «2Cl- + 2Hg(l) • Ecell = +0.24 V vs.SHE

  29. IR drop • Force needed to overcome resistance of ion movement • Follows Ohm’s law • Increase potential required to operate cell • ECell=Ecathode-Eanode-IR • For a Cd/Cu cell at 4 W, find potential needed for 0.1 A • Cu2+ + 2e- --> Cu 0.3394 • Cd2+ + 2e- --> Cd -0.4022 • Cu2++Cd<->Cu+Cd2+: • Ecell=0.3394-(-0.4022)-4*0.1=0.3416 V

  30. Polarization • ECell=Ecathode-Eanode-IR • predicts linear relationship between cell voltage and current • Deviation due to polarity of cell • Can occur at either electrode • Due to limitations of reaction at surface of electrode • Mass transfer • Concentration • Reaction intermediates • Physical processes • Sorption • Crystallization

  31. Methods