Lesson 4: Electrochemical Cells

• Electrochemical Cells


  Devices that convert chemical energy into electrical energy through a spontaneous redox reaction
• Standard Reduction Potential


  The tendency of a substance to undergo reduction (gain electrons)
• Alessandro Volta (1745-1827) invented the electric cell in 1800
• Voltaic cell


  A single electric cell, also called a galvanic cell or electrochemical cell
• Battery


  Multiple voltaic cells joined together
• Electrochemical cells


  • Use a spontaneous redox reaction to generate electrical energy by facilitating the passage of electrons through an external circuit
  • Convert chemical energy into electrical energy
• Voltage


  Potential difference
• Each electrochemical cell


  • Contains 2 electrodes (anode and cathode)
  • Contains electrolytes (aqueous conductors)
  • Contains a salt bridge (allows movement of ions)
  • Contains an external circuit (allows movement of electrons)
  • Contains an internal circuit (allows movement of ions)
• The electrodes can be a liquid or solution (rare case)
• Humans need electrolytes so there are no issues with the heart or brain, and humans are slightly salty
• Salt bridges replenish the electrons taken out of the cell
• Cathode


  Positive electrode
• Anode


  Negative electrode
• The salt bridge balances charges so electrons are still attracted to the positive side
• If there is a buildup of ions, the electrons will not be attracted, so the salt bridge keeps the electricity working
• Cell Notation


  A method of describing an electrochemical cell without drawing a diagram
• The salt bridge should not form a precipitate
• Group 1 metals never form a precipitate
• NO3 never forms a precipitate
• KNO3 is a common salt bridge material
• Electrochemical cell reaction


  1. Mg oxidized at anode

  2. Fe reduced at cathode

  3. Electrons flow from anode to cathode
• Electrochemical cell example 1

  • Mg (s) + Fe2+(aq) -> Mg2+(aq) + Fe(s)
• Cell Potential, Ecell


  The maximum voltage of the cell, depends on the composition of the electrodes and ion concentrations
• Standard Cell Potential, Eo

  cell


  The potential of the cell at standard conditions (1M, 25C, 1atm)
• The easier a substance undergoes reduction, the easier it attracts electrons
• The cell potential is maximized by the difference in reduction potentials of the two half-cells
• Reduction Potentials


  The tendency of a half-cell to undergo reduction (gain electrons)
• The half-cell with the larger reduction potential steals electrons from the half-cell with the lower reduction potential
• Standard Hydrogen Electrode (SHE)

  The reference electrode assigned a reduction potential of 0.00V
• The Standard Reduction Potential Table lists half-reactions in decreasing order of reduction potential
• Substances higher on the table are stronger oxidizing agents, substances lower are stronger reducing agents
• Calculating Standard Cell Potential

  1. Write oxidation and reduction half-reactions

  2. Obtain relevant reduction potentials

  3. Add the two values to get Eo

  cell
• Example 1: Calculating Eo

  cell for Ag-Cu cell

  • Eo

  cell = Eo

  cathode + Eo

  anode = 0.80V - 0.34V = 0.46V
• Example 2: Calculating Eo for Cu2+/H2 half-cell


  • Eo

  cell = 0.34V
• Example 3: Calculating cell reaction and Eo

  cell for Fe3+/Ni2+ cell


  • Cell reaction: 2Fe3+ + Ni -> 2Fe2+ + Ni2+

  Eo

  cell = Eo

  cathode + Eo

  anode = 0.77V + (-0.26V) = 1.03V
• Spontaneous reaction


  A reaction that occurs naturally without the addition of energy, indicated by a positive Eo

  cell
• Example 1: Determining spontaneity of Zn/Cr3+ reaction


  • Spontaneous, Eo

  cell = 0.76V + (-0.41V) = 0.35V
• Example 2: Determining spontaneity of Cu/H+ reaction


  • Nonspontaneous, Eo

  cell is negative
• Example 3: Determining spontaneity of Ag/MnO4- reaction

  • Spontaneous, oxidizing agent is higher on reduction potential table
• Non-spontaneous reactions can occur if electricity is added so the electrons move the other way (rechargeable batteries)