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

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

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

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+

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+

14-6 Periodic Variations of Oxidation State constant Steps of Steps of Mainly 3+

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

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

14-9 Balancing Redox Equations Balancing can be accomplished through examining ion- electron half reactions §H + + NO 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 e - §NO e - NO

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

14-11 Balancing Equations Add the reactions together §14H + + 2NO Cu 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+

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

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 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 e - Cu Anode §Electrode at which oxidation occurs §Anions migrate to anode à2Cl - Cl 2 + 2e -

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

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

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

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 C is the charge on 1 mole of electrons = 1F (faraday)

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

14-20 Electrochemical cell

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

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 e - Zn  °=-0.76 V §Fe 3+ is reduced, Zn is oxidized

14-23 Half-Cell Potentials Overall  2Fe 3+ +Zn 2Fe 2+ + Zn 2+  °= =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

14-24 Reduction Potentials Electrode Couple"E0, V" Na+ + e- --> Na Mg2+ + 2e- --> Mg Al3+ + 3e- --> Al Zn2+ + 2e- --> Zn Fe2+ + 2e- --> Fe Cd2+ + 2e- --> Cd Tl+ + e- --> Tl Sn2+ + 2e- --> Sn Pb2+ + 2e- --> Pb H+ + 2e- --> H2(SHE)0 S4O e- --> 2S2O Sn4+ + 2e- --> Sn SO H+ + 2e- --> H2O + H2SO3(aq) Cu2+ + e- --> Cu S + 2H+ + 2e- --> H2S AgCl + e- --> Ag + Cl Saturated Calomel (SCE) UO H+ + 2e- --> U4+ + 4H2O0.2682

14-25 Reduction Potentials Hg2Cl2 + 2e- --> 2Cl- + 2Hg0.268 Bi3+ + 3e- --> Bi0.286 Cu2+ + 2e- --> Cu Fe(CN)63- + e- --> Fe(CN) Cu+ + e- --> Cu0.518 I2 + 2e- --> 2I I3- + 2e- --> 3I H3AsO4(aq) + 2H+ + 2e- -->H3AsO3(aq) + H2O HgCl2 + 4H+ + 2e- -->Hg2Cl2 + 2Cl Hg2SO4 + 2e- --> 2Hg + SO I2(aq) + 2e- --> 2I O2 + 2H+ + 2e- --> H2O2(l) O2 + 2H+ + 2e- --> H2O2(aq) Fe3+ + e- --> Fe Hg e- --> Hg Ag+ + e- --> Ag Hg2+ + 2e- --> Hg Hg2+ + 2e- --> Hg NO3- + 3H+ + 2e- -->HNO2(aq) + H2O0.9275

14-26 Reduction Potentials VO2+ + 2H+ + e- --> VO2+ + H2O HNO2(aq) + H+ + e- --> NO + H2O Br2(l) + 2e- --> 2Br Br2(aq) + 2e- --> 2Br IO H+ + 10e- -->6H2O + I O2 + 4H+ + 4e- --> 2H2O MnO2 + 4H+ + 2e- -->Mn2+ + 2H2O Cl2 + 2e- --> 2Cl MnO4- + 8H+ + 5e- -->4H2O + Mn BrO H+ + 10e- -->6H2O + Br

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 = What is potential of an electrode of Zn(s) and 0.01 M Zn 2+ Zn 2+ +2e - Zn  °= V activity of metal is 1

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

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- --> Cu Cd2+ + 2e- --> Cd Cu 2+ +Cd Cu+Cd 2+: E cell = ( )-4*0.1= V

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

14-31 Methods