Unit 4.3.1: Electrode Potentials

3.1.11

Electrochemical cells

• Electrochemical cells can be made from two different metals dipped in salt solutions of their own ions (i.e. Fe (s) in Fe2+ (aq) ) and connected by a wire.

• There are always two reactions going on: Oxidation and Reduction - making this a redox process

  • The metal electrodes below are dipped in a s olution of their respective ions, connected via a salt bridge.
  • Zinc loses electrons more easily than copper so is by convention placed on the left.
  • This also means that the Zn electrode is more positive than the Cu electrode meaning the Zn(s) is oxidised to Zn2+, the electrons travel to the Cu2+ ions and reduces them to Cu(s)
  • The salt bridge allows ions to flow between the half-cells and balance the charges, completeing the circuit.

  

• Electrons flow through the wire from the most reactive to the least reactive.

• Cell Potential/Electromotive Force (EMF) - The voltage between two half-cells

You can also have a half-cell involving solutions of 2 aqueous ions of the same element

• A platinum electrode is used because it is inert (wont react)
• The conversion from Fe2+ to Fe3+ happens on the surface of the electrode

Electrode Potentials

• The reactions that occur at each electrode in a cell are reversible

  

• The direction the reaction goes in depends on how easily each metal is oxidised

• This is measured using electrode potential

• More negative electrode potential (V) = more easily oxidised

  • Zinc has a more -ve electrode potanital so is oxidised
  • Copper has a more +ve electrode potantial so is reduced

Drawing Electrochemical Cells

• Theres a shorthand way to drawing those big electrochemical cells:

  • | = state separation, you put this between substances of different phases in the same half-cell
  • || = salt bridge, you use this to separate the two half-cells

• Therefore, this electrochemical cell can be written as

  • Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu (s)

    

• By convention, unless it is Hydrogen, you always put the most negative electrode on the left.

• Follow these steps in order to draw electrochemical cells in shorthand:

  1. Use the electrode potentials to work out which haldf-cell goes on the left and which goes on the right (unless you have Hydrogen, which always goes on the left)
  2. Write out the left-hand half-equation as an oxidation reaction, and the right-hand half-equation as a reduction equation
  3. Write out the reactants and products of the oxidation reaction followed by the reactants and products of the reduction reaction
  4. Add in a salt bridge with || between the oxidation and reduction reactions
  5. Add in your phase separator, | , between any reagents that have different phases, and a comma between any of the same phase.
  6. If the element and electrode are separate, put the electrode on the outside and separate using a |

Calculating the cell potential

• The cell potential will always be a positive voltage because the more negative value is being subtracted from the more positive value