Redox and electrode potentials

Redox and electrode potentials

1. Redox reactions

   • A redox reaction is a reaction in which electrons are transferred
   • OIL RIG
     • Oxidation Is Loss (of e-)
     • Reduction Is Gain (of e-)
   • You can separate redox reactions into half equations
     • It is split into an oxidation half equation and a reduction half equation
     • You can only use H+, e- and H2O to balance half equations

2. Redox titrations

   • Redox titrations are used to find out how much oxidising agent is needed to react with a quantity of a reducing agent
   • Transition metals often change colour when they change oxidation state so they are useful in redox titrations
   • Titrations using transition element ions are redox titrations
   • To work out the concentration of a reducing agent you just need to know the known volume of an oxidising agent of a known concentration
   • Need to know examples

3. Iodine thiosulfate titrations

   • Iodine sodium thiosulfate reactions are a way of finding the concentration of an oxidising agent
   • The more concentrated an oxidising agent is the more ions will be oxidised by a certain volume of it
   • 1) Use a sample of oxidising agent to oxidise as much iodine as possible
   • 2) Find out how many moles of iodine have been produced
   • 3) Calculate the concentration of the oxidising agent

4. Electrochemical cells

   • Electrochemical cells make electricity
   • 1) Zinc loses e- more easily than copper so zinc is oxidised to form Zn2+ ions - this releases e- into the circuit
   • 2) In the other half cell, the same number of e- are taken from the external circuit, reducing copper ions to copper atoms
   • So the e- will flow through the wire from the most reactive metal to the least
   • Other Electrochemical cells
     • You can have half cells with solutions of the same element eg Fe2+ and Fe3+ - the conversion from Fe2+ to Fe3+ happens at the surface of the platinum electrode
     • You can also have electrochemical cells made from non metals - eg a gas can be bubbled over a platinum electrode sitting in a solution of aqueous ions
   • Electrode potential
     • The reactions that occur at each electrode is reversible
     • Which direction the reaction goes in depends on how easily each metal loses e- (oxidised)
     • Electrode potential is a measure of how easily a metal is oxidised - a metal that is easily oxidised has a very negative electrode potential
     • E(cell) = E(less reactive) - E(more reactive)
     • Electrode potentials are measured against standard hydrogen electrodes
     • Standard conditions:
       • 298K
       • 101KPa
       • 1mol dm-3
     • The standard hydrogen electrode potential is 0V
   • Electrochemical series
     • An electrochemical series shows you whats reactive and whats not
     • The more reactive a metal is the more it wants to lose e- to form a positive ion - more reactive metals have more negative electrode potentials
     • The more reactive a non metal is the more it wants to gain e- to form a negative ion - more reactive non metals have more positive electrode potential values
     • Changing any of the standard conditions can change the electrode potential due to the reactions being equilibrium reactions

5. Storage and fuel cells

   • Energy storage cells (battery) are like electrochemical cells
   • Fuel cells generate electricity from reacting a fuel with an oxidant - usually H with O
   • Advantages of electrochemical cells
     • More efficient at producing energy than conventional combustion engines - as energy in combustion is lost as heat
     • They produce a lot less pollution
   • Drawbacks of electrochemical cells
     • The production of the cells involves the use of toxic chemicals
     • The chemicals used to make the cell are often very flammable

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