3.1.11 electrode potentials

• What is a half cell?

  When a piece of metal id placed in a solution of its metal ions, and equilibrium is reached
• electrochemical cell


  when 2 electrodes are connected, potential difference between them creates a current, an electromotive force is created and the system can generate an electrical energy. Electrons flow through the wire from the more reactive metal to the less reactive one

  salt bridge- usually filter paper soaked in solution of KNO3
• metal electrodes

  metal surrounded by a solution of its ions
• gas electrodes

  gas and a solution of its ions, inert metal is the electrode to allow the flow of electrons
• redox electrodes

  2 different types of the same element where the 2 types of ions are present in solution with an inert metal electrode allowing the flow of electrodes
• measuring potential of an electrode

  Can only be measured between 2 electrodes, so it must be connected to another half-cell of known potential, and the potential difference between the 2 half cells is measured
• standard hydrogen electrode

  Primary standard that is used and is assigned the potential of 0 volts of a reference

  H is bubbled into a solution of H+ H+(aq) + e ↔  ½ H2(g)

  Standard conditions- 1 moldm-3 HCl, 100kPa, and 298K

  Components- 1 moldm-3 HCl, 100kPa, and Pt electrode
• electrochemical series

  Electrode potential measured using SHE are displayed in a table

  LHS= oxidising agents with the strongest at the bottom

  RHS= reducing agents with the strongest at the top
• Predicting reagents- answer framework

  The E$\theta$ for the __/__ half equation is more positive than the E$\theta$ for the __/__ half equation. Therefore __ is reduced to __ and __ is oxidised to __.
• E$^{\theta}$ cell


  reduction - oxidation
• conventional cell diagram

  Zn (s) | Zn 2+ (aq) || Cu2+ (aq) | Cu(s)

  | = phase boundary, a comma is used if cell components are in the same phase, or are a liquid and aqueous

  || = salt bridge
• conventional cell diagran of SHE

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• non-rechargeable fuel-cells

  • Chemicals are used up over time and EMF cells, once 1+ chemicals are used up, the cell is flat and EMF=0
  • Cannot be recharge and must be discharged off after their single use
  • Daniel cells= Zn/Cu cell, Zn runs out so cell cannot be recharged, solutions make it impractical for transport
  • Zn/C cells= Zn= negative electrode, C= positive electrode
  • Alkali batteries= higher cost, longer life
• Rechargeable fuel cells

  • Current (electrons) can be forced back the opposite direction by an applied external current
  • Half cell equations are reversable
  • Electrodes are designed so that they do not decompose or wear down as quickly. These can be very small, light and portable.
  • Lithium ion batteriesà electrolyte phase (liquid) is a polymer; this means the battery won’t leak
• Hydrogen fuel cells


  • continuous supply of the chemicals into the cell and so neither run out of chemicals nor need re-charging
  • two platinum electrodes separated by a special polymer electrolyte which will allow ions to pass through it
  • can be run in alkaline or acidic conditions, but the overall equation and the overall emf is the same
  • fuel cells don’t need recharging and EMF remains constant as there’s a continuous supply of fuel and concentrations remain constant
• using electrochemical series to calculate standard cell potentials

  E cell = E reduced - E oxidised

  (positive - negative)