Electrode potentials

• What you need to know about electrode potential

  • IUPAC convention for writing half equations for electrode reactions
  • Conventional representation of cells
  • Electrode potential and the standard hydrogen electrode
  • Importance of conditions when measuring electrode potential
  • Standard electrode potential
  • Electrochemical series
  • Using electrode potential values to predict direction of simple redox reactions
  • Calculating the EMF of a cell
  • Writing and applying the conventional representation of a cell
• Drawing a typical electrochemical cell

  1. Fill beaker with solution

  2. Add salt bridge

  3. Add electrodes

  4. Connect electrodes with wire and voltmeter
• Salt bridge


  Allows the flow of ions, not electrons
• Voltmeter


  High resistance to prevent current flow and measure maximum potential difference
• Electrode


  Solid metal
• Solution


  Contains metal ions of the same metal as the electrode
• Salt bridge should not be made of metal wire as it would set up its own electrode system
• Salt bridge should be made of an unreactive aqueous solution like saturated potassium nitrate to avoid reactions with the solutions
• Oxidation


  Loss of electrons
• Reduction


  Gain of electrons
• NOPR rule


  More negative half cell is oxidized, more positive half cell is reduced
• Comparing copper and zinc half cells

  Zinc is more negative, so it is oxidized. Copper is more positive, so it is reduced.
• Copper


  Positive 0.34 volts
• Zinc


  Negative 0.76 volts
• Copper and zinc can both be positive
• When one half-cell is more negative

  It is oxidized
• When one half-cell is more positive

  It is reduced
• Oxidation of zinc


  Zinc solid -> Zinc 2+ + 2e-
• Reduction of copper


  Cu 2+ + 2e- -> Cu solid
• Electrons produced in the oxidation reaction flow through the wire to the reduction reaction
• Conventional cell diagram


  Represents the two half-cells with a salt bridge in between
• The more positive half-cell is placed on the right side by convention
• If there is no solid metal electrode, a platinum electrode is used instead
• Electrode potential of one half-cell cannot be measured alone, it must be measured relative to another half-cell of known potential
• Standard hydrogen electrode


  Reference electrode with a potential of 0 volts
• Hydrogen gas is bubbled into the standard hydrogen electrode solution
• There's no such thing as solid hydrogen right we're not going to get a solid hydrogen electrode so what electrode do you think we're going to use for this Platinum all right we're going to use platinum electrode
• Drawing the half cell

  1. Draw electrode dipped in solution

  2. Connect to voltmeter and another half cell
• Drawing the standard hydrogen electrode

  1. Draw half cell

  2. Show hydrogen gas being bubbled in
• Platinum electrode


  It's not going to react with the hydrogen or the acid and it conducts electricity
• Reaction at the platinum electrode


  H2 gas ⇌ 2H+ (aq) + 2e-
• Standard conditions for standard hydrogen electrode

  • H2 gas at 100 kPa pressure
  • Concentration of H+ ions is 1 mol/dm3
  • Temperature is 298 K
  • Platinum electrode
• The standard hydrogen electrode has a potential of 0 V by definition
• Drawing the conventional cell diagram for standard hydrogen electrode

  H+ (aq) | H2 (g) | Pt (s)
• If the solution contains a strong acid like sulfuric acid, the concentration of H+ ions will be different from 1 mol/dm3
• Standard conditions


  • Constant concentration of all ions (1 mol/dm3)
  • Gases at 100 kPa pressure
  • Constant temperature (298 K)
  • No current flowing
• If the electrode half-reactions involve Fe2+ and Fe3+, their concentrations will both be 1 mol/dm3 under standard conditions
• Electrochemical series


  • Tabulated data showing reduction potentials of half-reactions
  • Reactions are written in the forward (reduction) direction
• Oxidizing agent


  Oxidizes something and gets reduced in the process
• Reducing agent


  Reduces something and gets oxidized in the process