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Electroanalytical chemistry

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

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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

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