Electro

• Electrochemistry


  The subdiscipline of chemistry that deals with the study of the relationship between electrical energy and chemical changes
• Electrochemical reactions


  • Chemical reactions that involve the input or generation of electric currents
• Categories of electrochemical reactions

  • Production of chemical change by electrical energy (electrolysis)
  • Conversion of chemical energy into electrical energy (generation of electricity by spontaneous redox reactions)
• Electricity can be produced when electrons move from one element to another in certain types of reactions (such as redox reactions)
• Electrochemistry deals with the overall reactions when multiple redox reactions occur simultaneously, connected via some external electric current and a suitable electrolyte

  
• Electrochemistry is also concerned with chemical phenomena that involve charge separation (as seen commonly in liquids such as solutions)
• The dissociation of charge often involves charge transfer that occurs homogeneously or heterogeneously between different chemical species
• Electrochemical cell


  A device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions
• Types of electrochemical cells


  • Galvanic cells
  • Electrolytic cells
• Galvanic cell


  Converts chemical energy into electrical energy (electricity can be obtained with the help of redox reaction)
• Example of galvanic cell


  • Daniell cell (zinc and copper used for a redox reaction)

  
• Redox reactions


  1. Oxidation half-reaction

  2. Reduction half-reaction

  
• Oxidizing agent


  Substance getting reduced in a chemical reaction

  
• Reducing agent


  Substance getting oxidized in a chemical reaction
• Examples of redox reactions


  • Reaction between hydrogen and fluorine
  • Reaction between zinc and copper
  • Reaction between iron and hydrogen peroxide

  
• Balancing redox reactions


  1. Using the change in oxidation number of oxidizing agent and reducing agent

  2. Dividing the redox reaction into two half reactions (one of reduction and one of oxidation)

  
• Examples of balancing redox reactions


  • Ferrous sulphate treated with acidified potassium permanganate
  • Reaction in acid medium
  • Reaction in basic medium

  
• SO4−2


  Sulfate ion
• H+


  Hydrogen ion
• SO3−2+H2O→SO4−2+2H++2e


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3

  
• 3SO3−2+3H2O→3SO4−2+6H++6e−


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3
• Cr2O7−2+14H++6e−→2Cr+3+7H2O


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3

  
• Cr2O7−2+3SO3−2→2Cr+3+3SO4−2+4H2O+8H+


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3

  4. Reaction step 4

  
• P4+12H2O+12e−→4PH3+12OHΘ


  Reaction step 1

  
• P4+8H2O+8OH−→4H2PO2−+4e−+8H2O


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3

  
• 3P4+24H2O+24OH−→12H2PO2−+12e−+24H2O


  1. Reaction step 1

  2. Reaction step 2

  3. Reaction step 3

  
• A spontaneous chemical process is the one which can take place on its own and in such a process Gibbs free energy of a system decreases
• In electrochemistry, spontaneous reaction (redox reaction) results in the conversion of chemical energy into electrical energy
• The reverse process is also possible where a non-spontaneous chemical reaction occurs by supplying electricity
• Galvanic cell


  • Converts chemical energy into electrical energy
  • Oxidation and reduction take place in two separate compartments
  • Each compartment consists of an electrolyte solution and metallic conductor which acts as an electrode

  
• Zn(s) + Cu+2(aq) → Zn+2 + Cu(s)

  1. Reaction step 1

  2. Reaction step 2
• Zn(s) →Zn+2 + 2e–

  Reaction step 1
• Cu+2(aq) + 2e– → Cu(s)


  Reaction step 1
• Salt bridge


  A device used in an electrochemical cell for connecting its oxidation and reduction half cells wherein a weak electrolyte is used
• Salt bridge


  • Helps maintain the electrical neutrality within the internal circuit
  • Prevents the cell from taking its reaction to equilibrium
  • Prevents the accumulation of positive and negative charges around the respective electrodes
  • Allows a smooth reaction to take place
  • Helps in the continual flow of electrons
• Electrode potential


  The tendency of losing of electrons (oxidation) or gaining of electrons (reduction)

  
• Standard electrode potential (E0)


  The electrode potential of an electrode relative to standard hydrogen electrode under standard conditions
• Electrochemical series


  A list that describes the arrangement of elements in order of their increasing electrode potential values
• Electrochemical series


  • Electrodes (metals and non-metals) in contact with their ions are arranged on the basis of the values of their standard reduction or oxidation potentials
  • Standard electrode potential is obtained by measuring the voltage when the half cell is connected to the standard hydrogen electrode under standard conditions

  
• Electropositive


  Elements (other than hydrogen) that display a greater tendency to lose electrons to their solution