electroanalytical chemistry n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Electroanalytical chemistry PowerPoint Presentation
Download Presentation
Electroanalytical chemistry

Loading in 2 Seconds...

play fullscreen
1 / 31

Electroanalytical chemistry - PowerPoint PPT Presentation


  • 222 Views
  • Uploaded on

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

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

PowerPoint Slideshow about 'Electroanalytical chemistry' - adamdaniel


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
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
  • Activities rather than concentrations
electroanalytical methods
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
oxidation and reduction
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
oxidation state
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+
oxidation state1
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+
periodic variations of oxidation state
Periodic Variations of Oxidation State

1

18

13-17

2

6-12

3

4

5

Steps of 1

constant

Steps of 2

Mainly 3+

oxidizing and reducing agents
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

redox reactions
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
balancing redox equations
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
balancing redox equations1
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)
balancing equations
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+
electrochemistry
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
electrochemistry1
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

electrolysis
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-
electrolysis1
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
electrolysis of cucl 2
Electrolysis of CuCl2

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

Cathode: Cu2++2e-->Cu

C electrode

C electrode

Cu Plating on C electrode

Cl2

nacl solutions
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
faraday laws
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)
faraday laws1
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
conduction in a cell
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
half cell potentials
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
half cell potentials1
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
reduction potentials
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

reduction potentials1
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

reduction potentials2
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

nernst equation
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
electrodes
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
ir drop
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
polarization
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