1.11 Electrode potential and chemical cells

• Dipping a metal rod into a solution of its own ions


  An equilibrium is set up between the solid metal and the aqueous metal ions
• Zinc (s) to zinc (II)


  Zn (s) ⇌ Zn2+(aq) + 2e-
• Copper (II) to copper (III)


  Cu2+(aq) ⇌ Cu3+(aq) + e-
• Salt bridge


  Filter paper soaked in saturated solution of KNO3 (potassium nitrate)
• Purpose of salt bridges

  Complete the circuit, but avoid further metal/ion potentials as does not perform electrochemistry. Allows ion movement to balance the charge. Do not react with electrodes
• Symbol for salt bridge in standard notation


  ||
• Species placed on the outside (furthest from the salt bridge) in standard cell notation

  The most reduced species
• What | indicates


  Phase boundary (solid/liquid/gas)
• Aluminium/Copper cell representation


  • Al(s) | Al3+(aq) || Cu2+(aq) | Cu(s)
• Left-hand electrode


  Oxidation occurs, has the most negative Eo value
• Right-hand electrode


  Reduction occurs, has the most positive Eo value
• The side with the most negative Eo value undergoes oxidation
• Conditions for standard hydrogen electrode

  Temperature = 298 K, Pressure = 100 kPa, [H+] = 1.00 mol dm-3
• Use of standard hydrogen electrode


  Comparing other cells against. EO of SHE is defined as 0, so all other Eo values are compared against it.
• Reasons for using other standard electrodes

  They are cheaper/easier/quicker to use and can provide just as good a reference. Platinum is expensive.
• Meaning of more negative Eo value

  Better reducing agent (easier to oxidise)
• Meaning of more positive Eo value

  Better oxidising agent (easier to reduce)
• Factors that change Eo values

  • Concentration of ions, Temperature
• Reducing concentration of ions in left-hand half cell

  Equilibrium moves to the left to oppose the change of removing ions; this releases more electrons, the Eo of the left hand cell becomes more negative, so the e.m.f. Of the cell increases.
• Calculating cell emf from Eo values

  Eo

  cell = Eo

  right - Eo

  left
• When to use a Platinum electrode

  When both the oxidised and reduced forms of the metal are in aqueous solution
• Why Platinum is chosen

  Inert so does not take part in the electrochemistry, Good conductor to complete circuit
• Predicting if a reaction will occur

  Take the 2 half equations, find the species that is being reduced (this is effectively the right hand electrode), calculate its Eo value minus the Eo value of the species that is being oxidised (effectively the left hand electrode)
• Platinum electrode


  Used when both the oxidised and reduced forms of the metal are in aqueous solution
• Platinum is chosen because it is inert and does not take part in the electrochemistry, and it is a good conductor to complete the circuit
• Predicting if a reaction will occur

  1. Take the 2 half equations

  2. Find the species that is being reduced (this is effectively the right hand electrode)

  3. Calculate its Eo value minus the Eo value of the species that is being oxidised (effectively the left hand cell)

  4. If Eo overall > 0, reaction will occur
• The first commercial cell (Daniell cell) was made from zinc/copper(II)
• Zinc/carbon cells


  Disposable batteries
• Reactions in zinc/carbon cells

  1. Zn oxidised to Zn2+

  2. NH4+ reduced to NH3 at carbon electrode
• Reactions in lead/acid batteries (car batteries)

  1. Pb + SO42- → PbSO4 (s) + 2e-

  2. PbO2 + 4H+ + SO42- + 2e- → PbSO4 + 2H2O
• Recharging rechargeable cells


  Reactions are reversible and are reversed by running a higher voltage through the cell than the cell's Eo
• Reactions in nickel/cadmium rechargeable cells

  1. Cd(OH)2 (s) + 2e- → Cd(s) + 2OH-

  2. NiO(OH) (s) + H2O + e- → Ni(OH)2 (s) + OH-
• Lithium-ion cells are used in mobile phones and laptops
• Reactions on discharge in lithium-ion cells

  1. Li+ + CoO2 + e- → Li+[CoO2]-

  2. Li → Li+ + e-
• Fuel cell


  A cell that is used to generate electric current; does not require electrical recharging
• Reactions in an alkaline hydrogen fuel cell


  1. 2H2 + 4OH- → 4H2O + 4e-

  2. O2 + 2H2O + 4e- → 4OH-
• Diagram of a hydrogen fuel cell
• Using a fuel cell vs burning H2 in air

  Fuel cell is more efficient and does not produce sulfur and nitrogen containing compounds that are bad for the environment
• Disadvantages of fuel cells

  • Hydrogen is a flammable gas with a low b.p. → hard and dangerous to store and transport → expensive to buy
  • Fuel cells have a limited lifetime and use toxic chemicals in their manufacture
• Finding the weakest reducing agent from a table of electrode potential data

  Most positive Eθ value. Then it is the PRODUCT of the reduction equation i.e. imagine equation going from right to left