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Electrochemistry

• Electrochemistry is the chemistry of the conversion of chemical energy to electrical energy through the exchange of electrons
  
  
• Similarly an electric energy can be converted into chemical energy.
  
  Related topics -
  
  
  • Electrode
    
    
  • Redox Potentials
    
    
  • Electrochemical Reaction
    
    
  • Electrochemical Cell
    
    
  • Half Cell
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
• The Electrode
  
  
  • A simple electrode can be a metal wire, along which an electric current can flow (ie a conductor of electricity)
    
    
  • Typically, electrodes come in pairs connected by a wire and dipped into a solution to form an electric circuit (see diagram)
    
    
  • Electrodes can be composed of different materials, usually metals, with each metal able to release electrons with varying degrees of ease. (ie redox potentials).
    
    
    
• Redox Potentials
  
  
  • Zinc (Zn) or a Zn (zinc) electrode will release electrons more readily then then copper (Cu) or a Cu (copper) electrode, and a silver (Ag) electrode will release electrons more readily then either Zn or Cu electrodes.
    
    
  • As an example, the loss of electrons from an atom of Zn is referred to as oxidation.
    
    
    Zn => Zn⁺⁺ + 2 e^- (e^- represents a released electron)
    
    
    i.e. A zinc atom (Zn) becomes, or is oxidized to a Zn ion (Zn⁺⁺) and two electrons (2e^-).
    
    
  • Because the electrons (e^-) can not exist in a free state they will have to be accepted by another substance or oxidant with a greater affinity for the electrons.
    
    
  • A Cu⁺⁺ (copper ion), acting as a oxidant, accepts the two electrons and is itself reduced to elemental Cu.
    
    
    Cu⁺⁺ + 2e^- => Cu
    
    
    i.e. A copper ion (Cu⁺⁺) on receiving two electrons (2e^-) becomes, or is reduced to a elemental copper (Cu).
    
    
  • That is, reduction and oxidation always occur together or are coupled as redox reactions.
    
    
  • A redox reaction can be seen as a tug-of-war for electrons with winners and loosers.
    
    
  • The tendency of a substance to release electrons or the strength of the electron bond is called the redox potential and is measured in volts (V).
    
    All redox potentials are measured relative to a hydrogen electrode which is set at an arbitrary 0.00 V (volts).
    
    The electrical or redox potential between the two following half reactions are -
    
    
    Zn => Zn⁺⁺ + 2 e^- and 2H⁺+ 2 e^- =>H₂ is +0.76 volts.
    
    Cu => Cu⁺⁺ + 2 e^- and 2H⁺+ 2 e^- =>H₂ is -0.34 volts.
    
    Or
    
    H₂ =>H₂ +2H⁺+ 2 e^- and Cu⁺⁺ + 2 e^- => Cu is +0.34 volts.
    
    
    This means the driving force or redox potential between the paired redox half reactions below can be calculated by each half reaction's redox potential -
    
    
    Zn => Zn⁺⁺ + 2 e^- +0.76 volts.
    
    Cu⁺⁺ + 2 e^- => Cu +0.34 volts.
    
    Total = + 1.10volts.
    
    
    
  • Electrochemical Reaction
    
    
    • Copper sulphate (CuSO₄) in solution ionizes to copper ions (Cu⁺⁺) and sulphate ions (SO₄⁼.
      
      
    • A piece of Zn metal placed in a CuSO₄ solution will give up two electrons per atom to Cu⁺⁺ producing Zn⁺⁺ and elemental Cu (as shown in the above equations). The Cu can gradually be observed as a brown copper deposit on the Zn metal and the Zn metal will gradually be consumed.
      
      
    • A piece of Cu metal placed in a ZnSO₄ solution will give up two electrons per atom to Zn⁺⁺ producing elemental Zn and Cu⁺⁺ ions.
      
      
      
  • Electrochemical Cell
    
    
    • As discussed above, electrochemical reactions involve a flow of electrons to a substance that have a greater affinity for them.
      
      
    • The flow of these electrons can be harnessed as an electric current as supplied by batteries or the electric potential used to measure a chemical environment (as in a pH meter).
      
      
    • The electrochemical cell, as set up as in diagram 1, will result in a unharnessed, direct electron flow within the solution.
      
      
    • The electrochemical cell, as set up in diagram 2, will result in a harnessed electron flow through the wire connecting the electrodes.
      
      
    • The electrochemical cell, as set up in diagram 3, will not result in an electrical current as the circuit is incomplete.
      
      
    • The salt bridge completes the electric circuit.
      
      The salt bridge can consist of a porous membrane between solutions, a paper or porous wick or a bridge containing a gel of a suitable salt (eg KNO₃).
      
      The KNO₃ exists in ionic form to replace the Cu⁺⁺ as they form elemental Cu. This occurs and needs to occur to maintain the charge (+ and -) balance.
      
      Remove the salt bridge and no current will flow.
      
      
      
  • Half Cell
    
    
    • An Electrochemical half cell is an old term used to describe one of the separated units of an electrochemical cell, consisting of a electrode and the surrounding electrolyte (ionic solution) - see diagram 4
      
      
    • The reference electrode and the glass electrode of a pH meter can both be regarded as half cells.
      
      
    • The reference electrode is required to provide a stable electrical reference potential against which the pH measuring electrode electrical potential can be ascertained
      
      
      
  • Reference Potential
    
    
    • The reference electrode consists of an internal silver/silver chloride (Ag/AgCl) electrode/wire and an internal, pH neutral (pH 7) solution of concentrated potassium chloride (KCl)
      
      
    • This combination half cell provides a stable reference potential.
      
      
    • This combination half cell provides an electric potential of 0.205 volts (V) (@ 25^oC. vs.NHE, 3.5 M KCl).