Redox Equilibria

Redox Equilibria

1. Redox Equations

   • oxidation is loss
   • Reduction is Gain
   • be able to write half equations
     • add water to balance hydrogen

2. Electrode potentials

   • IUPAC representation of cells
     • Use I for a phase difference and II for salt bridge
     • Be very careful with state symbols, use platinum when there is no solid electrode
     • ROOR, reduced I oxidised II oxidised I reduced
       • "negative people go on the outside"
   • Standard Conditions
     • 1 M concentration
     • 1 bar/ 100kpa
     • 298K

3. Electrochemical Series

   • electrode potentials can be listed as a series, +ve to -ve or visa versa
   • the more negative a potential is, the better the electrode is at donating electrodes, so it is a better reducing agent and is itself oxidised
   • emf of a cell =RHS-LHS
   • effect of condition on cell emf
     • apply Le Chatelier's principle

4. Standard Hydrogen Electrode

   • Composed of Hydrogen gas bubbled into a solution of H+ ions and platinum electrode
     • platinum electrode: inert metal used as either a source or sink of electrons
   • Can't measure the potential of single electrode, so SHE is assigned "0v" to allow comparison of other electrodes
   • others may be used because hydrogen is explosive and often unreliable/complex to set up

5. Electrochemical Cells

   • Primary Cells
     • Irreversable, provide current to external circuit whilst discharging and are disposed after
   • Secondary Cells
     • Designed to be recharged by an external circuit, lead acid is an example
   • Fuel Cell
     • generates electricity from the continuous oxidation of an external fuel source(hydrogen)
     • overall reaction is the exact same as combustion, but it is much more efficient since it converts directly to electrical energy
     • looks clean initially but most hydrogen is produced using non renewable resources which leave a large carbon footprint
       • also can be dangerous since hydrogen is highly flammable