Electrochemical cells

• A half cell is when a rod of metal (electrode) is dipped into a solution of its own ions, an equilibrium is set up
• Electrochemical cell

  Formed when two half cells are joined together with a wire, a high resistance voltmeter and a salt bridge

  • It's a redox process
  • One half cell is oxidised and the other is reduced
  • The more reactive metal is (the one that looses electrons more easily ) the half cell which undergoes oxidation
  • The half cell that undergoes reduction contains the less reactive metal
• • The salt bridge is made of KNO3 and it allows ions to flow between the half cell and balance the charges
• Some half-cells don't contain solid metal so a platinum electrode is used as its inert but will conduct electricity
• Conditions required

  • 298k
  • All solutions at 1.00 moldm^-3
  • Any gases at 100kPa

  These will get the same values and can be used to compare cells
• An electrochemical cell can be represented in a shorthand way by a cell diagram called the conventional representation or standard cell diagram
• Standard cell diagram

  Label the diagram

  A) Phase change

  B) Phase change

  C) Salt bridge

  D) Oxidised

  E) Reduced

  F) mix

  
• The standard hydrogen electrode by definition has a e.m.f potential difference of 0v ( 0 volts)
• Label diagram

  A) H2 gas

  B) 100 kPa

  C) 298k

  D) left

  E) Platinum electrode

  
• If a substance is in the same phase e.g. Fe2+ and Fe3+ then when writing the cell diagram they would be separated using a comma
• Label the diagram

  A) Salt bridge

  B) Platinum

  C) |

  D) ||

  E) ,

  F) |

  G) Pt(s)

  
• The reaction with the most negative electrode potential is the better reducing agent ( so more likely to be oxidised)
• The best reducing agent is the one species that looses electrons easily
• Electrode potentials that are negative potential values of E°/v are better at releasing electrons (better reducing agents) than hydrogen
• Changing conditions (temp, pressure or conc)

  Would cause equilibrium to shift, producing more electrons causing the electrode potential to be more negative
• Ecell = E°(RHS) - E° (LHS)

  This is used to calculate the e.m.f
• For a reaction to be feasible the Ecell must be positive
• When combining two electrode potential equation, you swap the most negative one which would be the electrode that is on the LHS and wants to be oxidised