Redox equilibria -5


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  • Created by: Shannon
  • Created on: 02-03-15 11:17
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  • Redox equilibria
    • Redox potentials/Electrode potentials
      • Tell us the power or potential of an aqueous species to oxidise or reduce
        • A +ve E*/V will produce a spontaneous reaction
          • The more +ve the value, the more likely it is to happen
      • Occur under standard conditions
        • 298 K
        • 1 moldm-3
        • 1 atm
        • Hydrogen half cell is used as a reference point = 0.00 E*/V
          • Hydrogen electrode attached to the electrode system being investigated
            • Connected via a salt bridge
            • Always the left hand component of a cell
        • The plate making contact with the solution must be platinum if the reduced form of the system we are studying i not a conducting metal
      • E* values produce a way of comparing positions of equilibrium when elements lose electrons to form ions in solution
        • The more -ve the E*, the further to the left the equilibrium lies
          • More likely to lose electrons to form ions
        • The more +ve the E* the further to the right the equilibrium lies
          • Less likely to lose electrons to form ions
      • Why is a reaction not feasible is E*/V +ve?
        • Reaction takes a more favoured route than the 1 we are investigating
        • Activation energy is too high
        • Conditions are not standard
    • Cell short hand notations
      • Oxidation takes place at the ANODE
      • Reduction takes place at the CATHODE
      • Phase boundary represented by '/'
      • Salt bridge represented by '//'
      • If two species are in the same phase, they are separated using a comma, e.g Cu+,Cu2+
      • Conventionally, the hydrogen half cell is always written on the left
    • Reduction & Oxidation
      • Reduction is a gain of electrons
        • A reducing agent donates electrons to the oxidising agent it reduces
      • Oxidation is a loss of electrons
        • An oxidising agent accepts electrons from the reducing agent it oxidises
      • Reduction & oxidation can be represented by half equations
        • These equations can be combined to give a full equation
    • Hydrogen fuel cell
      • Cell set up - polymer electrolyte membrane between anode and cathode, only allows H+ to pass through, e- have to travel through circuit to get to cathode - creating an electric current
      • Advantages
        • Efficiency almost 100% - more efficient over normal car batteries
        • Only waste product is water - less pollution than normal car batteries
      • Disadvantages
        • Production of H2 requires more energy than the fuel cell provides
        • Expensive and limited life-span
    • Alcohol fuel cell
      • Either uses methanol or ethanol - as H2 rich fuels
        • Made from renewable biomass
      • Electrons transported through an external circuit from anode -> cathode
      • Advantages
        • Higher hydrogen density the H2 fuel cell
        • Don't need special refrigeration like H2 - as liquid at room temp
      • Disadvantages
        • Low efficiency as alcohols can pass through membranes
        • Methanol is toxic
    • Breathalysers
      • Modern breathalysers VS Police breathalysers VS Old breathlysers
        • Modern breathalysers are easily portable and accurate
          • Ethanol fuel cell - amount of alcohol in breath in proportional to the current produced when breath is fed to the anode of the cell
        • Police breathalysers use IR to detect the presence & quantity of ethanol
          • Accurate but not easily portable
        • Old breathlysers used the reduction of potassium dichromate

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