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 BatteriesProduction of metals from ores

• Oxidation -reduction occurs simultaneouslyFor atoms and monatomic ions, loss or gain

of electronsFor 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 stateNaCl: Na+

and Cl-

MnO4-: Mn7+ and 4O2-

• For free elements each element is assigned an oxidation state of 0HgCl2

P4

• For monotonic ion, oxidation state is the chargeCl-

, Pu4+

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)

NaH, LiHLiAlH4, NaBH4

• Group 2 (IIA) elements (Be, Mg, Ca, Sr, Ba, and Ra) are 2+• Oxygen is usually 2-

Exceptions with oxygen-oxygen bondsH2O2, Na2O2: O oxidation state = 1-KO2: O oxidation state = 1/2-OF2: O oxidation state = 2+

14-6

Periodic Variations of Oxidation State

constant

1

2

3

Steps of 1

4 5 6-12

Steps of 2

13-17

18

Mainly 3+

14-7

Oxidizing Agents Reducing AgentsF2 F-

Cl2 Cl-

Br2 Br-

Ag+ AgI2 I-

Cu2+ CuH+ H2

Fe2+ FeZn2+ ZnAl3+ AlNa+ Na

Oxidizing and Reducing Agents

weak

Strong

14-8

Redox Reactions

• Zn + Cu2+ <--> Zn2+ + CuZn is oxidized, Cu is reducedTransfer of electrons from one metal to another

• May not involved charge speciesC + O2 <--> CO2

• 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 reactionsH+ + NO3

- + Cu2O <--> Cu2+ + NO + H2O• Identify reduced and oxidized species

Cu2O to Cu2+ (1+ to 2+): oxidizedNO3

- to NO (5+ to 2+): reduced• Balance oxidized/reduced atoms

Cu2O <--> 2Cu2+

• Add electrons to balance redox of elementCu2O <--> 2Cu2+ + 2e-

NO3- + 3e- <--> NO

14-10

Balancing Redox Equations

• Add H+ (or OH-) to balance charge of reaction2H+ + Cu2O <--> 2Cu2+ + 2e-

4 H+ + NO3- + 3e- <--> NO

• Add water to balance O and H, then balance other atoms if needed2H+ + Cu2O <--> 2Cu2+ + 2e- + H2O

4 H+ + NO3- + 3e- <--> NO + 2 H2O

• Multiple equations to normalize electrons3(2H+ + Cu2O <--> 2Cu2+ + 2e- + H2O)

2(4 H+ + NO3- + 3e- <--> NO + 2 H2O)

14-11

Balancing Equations

• Add the reactions together14H+ + 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 agent2 Pu4+ <--> Pu3+ + Pu5+

14-12

Electrochemistry• Chemical transformations produced by

electricityCorrosionRefining

• Electrical UnitsCoulomb (C)

Charge on 6.25 x 1018 electronsAmperes (A)

Electric currentA=1C/sec

14-13

Electrochemistry• Volt (V)

Potential driving current flowV= 1 J/C

• Ohm’s law = IR

= potential, I =current, and R=resistance

symbol unit relationshipsCharge q Coulomb (C)Current I Ampere (A) I=q/t (t in s)Potential Volt (V) =IRPower P Watt (W) P= IEnergy E Joule (J) Pt= It= qResistance R Ohm () R= /I

14-14

Electrolysis

• Production of a chemical reaction by means of an electric current2 H2O <--> 2H2 + O2

• CathodeElectrode at which reduction occursCations migrate to cathode

Cu2+ + 2e- <--> Cu• Anode

Electrode at which oxidation occursAnions migrate to anode

2Cl- <-->Cl2 + 2e-

14-15

Electrolysis

• Redox depends upon tendencies of elements or compounds to gain or lose electronselectrochemical series

Lists of elements or compounds Half cell potentials

• Related to periodic tendencies

14-16

Electrolysis of CuCl2

C electrode

Cu Plating on C electrode

C electrode

Cl2

Anode: 2Cl-->Cl2+2e-

Cathode: Cu2++2e-->Cu

14-17

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 Saltanode: 2Cl- <--> Cl2 + 2e-

cathode: Na+ + e- <-->NaNa metal produced by electrolysis of NaCl and Na2CO3

Lower 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

• 96487 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 electrodeanode: Cu <--> Cu2+ + 2e-

cathode: Cu2+ + 2e- <--> CuNumber of electrons

(10A)(3600 sec)/(96487 C/mol) = 0.373 F0.373 mole e- (1 mole Cu/2 mole e-) =

0.186 mole Cu

14-20

Electrochemical cell

14-21

Conduction in a cell

• Charge is conducted Electrodes Ions in solution Electrode surfaces

Oxidation and reductionOxidation at anodeReduction at cathode

• Reaction can be written as half-cell potentials

14-22

Half-cell potentials• Standard potential

Defined as °=0.00VH2(atm) <--> 2 H+ (1.000M) + 2e-

• Cell reaction for Zn and Fe3+/2+ at 1.0 MWrite as reduction potentials

Fe3+ + e- <--> Fe2+ °=0.77 VZn2+ + 2e- <-->Zn °=-0.76 V

Fe3+ is reduced, Zn is oxidized

14-23

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

14-24

Reduction PotentialsElectrode Couple "E0, V"Na+ + e- --> Na -2.7144Mg2+ + 2e- --> Mg -2.3568Al3+ + 3e- --> Al -1.676Zn2+ + 2e- --> Zn -0.7621Fe2+ + 2e- --> Fe -0.4089Cd2+ + 2e- --> Cd -0.4022Tl+ + e- --> Tl -0.3358Sn2+ + 2e- --> Sn -0.141Pb2+ + 2e- --> Pb -0.12662H+ + 2e- --> H2(SHE) 0S4O62- + 2e- --> 2S2O32- 0.0238Sn4+ + 2e- --> Sn2+ 0.1539SO42- + 4H+ + 2e- --> H2O + H2SO3(aq) 0.1576Cu2+ + e- --> Cu+ 0.1607S + 2H+ + 2e- --> H2S 0.1739AgCl + e- --> Ag + Cl- 0.2221Saturated Calomel (SCE) 0.2412UO22+ + 4H+ + 2e- --> U4+ + 4H2O 0.2682

14-25

Reduction PotentialsHg2Cl2 + 2e- --> 2Cl- + 2Hg 0.268Bi3+ + 3e- --> Bi 0.286Cu2+ + 2e- --> Cu 0.3394Fe(CN)63- + e- --> Fe(CN)64- 0.3557Cu+ + e- --> Cu 0.518I2 + 2e- --> 2I- 0.5345I3- + 2e- --> 3I- 0.5354H3AsO4(aq) + 2H+ + 2e- -->H3AsO3(aq) + H2O 0.57482HgCl2 + 4H+ + 2e- -->Hg2Cl2 + 2Cl- 0.6011Hg2SO4 + 2e- --> 2Hg + SO42- 0.6152I2(aq) + 2e- --> 2I- 0.6195O2 + 2H+ + 2e- --> H2O2(l) 0.6237O2 + 2H+ + 2e- --> H2O2(aq) 0.6945Fe3+ + e- --> Fe2+ 0.769Hg22+ + 2e- --> Hg 0.7955Ag+ + e- --> Ag 0.7991Hg2+ + 2e- --> Hg 0.85192Hg2+ + 2e- --> Hg22+ 0.9083NO3- + 3H+ + 2e- -->HNO2(aq) + H2O 0.9275

14-26

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

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

Zn2+

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

2.30RT

nFlog

[C]c[D]d

[A]a[B]b

0.763 0.0592

2log

10.01

0.822V

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 Hg2Cl2 (s) + 2e- «2Cl- + 2Hg(l) Ecell = +0.24 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 ECell=Ecathode-Eanode-IR

• For a Cd/Cu cell at 4 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

14-30

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

14-31

Methods