april 30

electrochemistry

a few misc. topics

• galvanic vs electrolytic cells
  • batteries and pH meters are galvanic
  • spontaneous, voltage / current sources
• voltammetry, electrolysis, recharging a cell are electrolytic cells
  • external power supply needed
  • circuit is driven; not spontaneous
  • could be the reverse of a spontaneous process
• Faradaic vs. non-Faradaic Currents
  • Faraday's Law-- electrons go into redox reaction
    • where else could they go?
  • For short periods, current can build up charges
    • in solution, at electrode
    • so may see initial current surges (1 msec) that are not Faradaic
• Related Problem
  • if local charges are created around an electrode
  • then measurement is not of the original solution
  • (conductivity usually AC to avoid "polarizing" the solution)
• anode vs cathode
  • anode-- reaction where oxidation occurs (two vowels)
  • cathode-- reaction where reduction occurs
  • (rechargeable cell reverses anode/cathode)
• galvanic cell
  • oxidation (Cu-->Cu²⁺ +2e) is electron source
  • therefore anode is negative
• electrolytic cell
  • reduction (Cu²⁺ + 2e- --> Cu) needs electrons
  • cathode must be negative (excess electrons)
• Double Layer
  • electrode -- adjacent solution -- rest of solution
  • the first 5-100 microns around electrode has composition and change unlike either the electrode or the bulk solution
• As a rule, a few milliamps causes no significant concentration change beyond electrode
  • 1 mole of chemistry = 96,500 amp-sec
  • 1 ma would need 96,000,000 seconds to cause a mole of reaction
    • that's 3 years
  • so, a pH meter lives by reaction of H⁺, but need not worry about changing the concentration by the measurement process.
    • (worry a bit if concentrations are 10^-9 molar or volume is tiny)
    • (for a battery with 0.1 mole of material, 1 ma for 3 months)
  • even the double layer effect is negligible for potentiometric determinations
    • can be a concern when current flows (voltammetry, conductivity)
• Reference Electrodes
  • to act as the second electrode (anode or cathode)
  • to provide same voltage no matter what's in the solution
  • examples (figures in text)
    • calomel-- Hg(s) | Hg₂Cl₂ (s) | Cl-
    • most convenient if Cl- = saturated KCl
    • a small evaporation keeps concentration fixed
    • solubility is only slightly temperature dependent
    • will be saturated with Hg₂Cl₂ also
  • small junction into the test solution
    • crack in glass, fiber, polymeric material, porpus ceramic, glass frit, porous glass
      • will be a small (variable) liquid junction potential
      • difficult to determine for precise work
  • Ag|AgCl(s)| Cl-
    • silver wire, thin coating of AgCl
      • simpler than calomel; gradually taking over
      • historically, too much data is referenced to calomel
    • coating will last "forever" at these currents
    • typically saturated KCl
      • also other Cl- conc (Normal, deci-normal)
    • May add a KNO₃ sleeve if Cl- is a problem
    • leakage is very slow (<1 ml/year)
    • but Cl- crust can be a Cl- source
• Ideal electrode
  • reaction is rapid so always at equilibrium with immedaite surroundings
  • reaction is specific-- only one solute reacts, producing a potential
  • electrode isn't poisoned or damaged by other species
  • Faraday's Law applies (to concentration) without any approximations
  • electrical resistance is moderate (1000 ohm, for example)
  • set up is simple (no H₂ tanks)
• Redox Electrodes
  • typically Pt (Au, graphite)
  • electrode itself is not reactive; just a connector
    • facilitates oxidation or reduction
    • however, surface plays a role-- perhaps catalytic
    • so may be a real difference between Pt, Au, C, SnO₂
  • measures the potential associated with any (all) redox reactions
  • after all, if there were several processes not at equilibrium
  • then one redox reaction would occur, reacting with the other system
  • usually we only have one system under study
  • consider a titration of Fe²⁺ with redox reagent like Cr₂O₇2- or Ce⁺⁴
  • once started we have Cr₂O₇2- and Cr³⁺ and a redox potential
  • but we might have Fe³⁺ and Fe²⁺ and a redox potential for these
  • prior to the endpoint Fe3+/Fe2+ is a modest ratio
  • easily used in the Nernst Equation
  • the Cr₂O₇2- doesn't remain, so it is extremely tiny
  • Nernst equation with Cr₂O₇2- / (Cr³⁺)² is correct, but outrageous ratios
  • after the equivalence point it's easier to work with the Cr₂O₇2- / (Cr³⁺)² ratio
  • at the equilence point, both equations are probably needed