Lesson 1: Intro to Electrochemistry

• Electrochemistry

  The study of all reactions in which a transfer of electrons takes place
• Types of electrochemical reactions

  • Some reactions produce electricity e.g., batteries
  • Some reactions require electricity e.g., electrolysis of water
  • Some occur where no electricity is produced or used redox reactions are involved in electrochemistry
• Oxidation

  The loss of one or more electrons
• Reducing agent

  The substance being oxidized (losing electrons)
• Reduction

  The gain of one or more electrons
• Oxidizing agent

  The substance being reduced (gaining electrons)
• OIL

  Oxidation is lost
• RIG

  Reduction is gain
• Reduction refers to overall charge, so the overall charge is reduced when it gains electrons</b>
• It is a reducing agent because it causes the other substance to get reduced
• It is an oxidizing agent because it causes the other substance to get oxidized
• Nonionic Equation


  Zn(s) + CuSO4(aq) → Cu(s) + ZnSO4(aq)
• Total Ionic Equation


  Zn(s) + Cu2+(aq) + SO42-(aq) → Cu(s) + Zn2+(aq) + SO42-(aq)
• Net Ionic Equation


  Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq)
• The blue colour of the CuSO4(aq) fades as Cu(s) is deposited on the Zn(s), which dissolves in solution to form ZnSO4(aq)
• In terms of electron transfer, Cu2+ gains 2 electrons and Zn(s) loses 2 electrons to Cu2+
• There are 2 half reactions that together form the overall redox equation
• As the reaction proceeds, the blue color of CuSO4 fades and is deposited to Zn to become a colourless ZnSO4
• Cu2+ gains electrons while Zn loses 2 electrons
• Only aqueous substances can dissociate into its separate ions
• Half Reactions


  1. Zn(s) → Zn2+(aq) + 2 e- (Zn loses electrons, Oxidation half reaction)

  2. Cu2+(aq) + 2 e- → Cu(s) (Cu2+ gains electrons, Reduction half reaction)
• Overall


  Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
• The reaction is balanced for the number of atoms/ions and charge
• Gain of electrons means it is a reactant
• Loss of electrons means it is a product
• Zn causes copper to become reduced by transferring electrons
• Nonionic Equation


  Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
• Net Ionic Equation


  Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
• Half Reactions


  1. Cu(s) → Cu2+(aq) + 2e-

  2. 2Ag+(aq) + 2e- → 2Ag(s)
• Cu is oxidized, it is the reducing agent
• Ag+ is reduced, it is the oxidizing agent
• Nonionic Equation


  2Na(s) + Cl2(g) → 2NaCl(s)
• Net Ionic Equation


  2Na(s) → 2Na+(aq) + 2Cl-(aq)
• Half Reactions


  1. 2Na(s) → 2Na+(aq) + 2e-

  2. Cl2(g) + 2e- → 2Cl-(aq)
• Na is oxidized (reducing agent)
• Cl2 is reduced (oxidizing agent)
• Total Ionic Equation


  Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)
• Half Reactions


  1. Zn(s) → Zn2+(aq) + 2e-

  2. 2H+(aq) + 2e- → H2(g)
• Balancing Redox Reactions


  • Write the half reactions (including number of electrons)
  • Balance each half reaction for number of atoms/ions
  • Balance each half reaction for charge: multiply one or both of the half reactions by a value such that the number of electrons lost during oxidation must equal the number of electrons gained during reduction
  • Add the two half reactions to give the overall equation
• Example 1


  • Ag+(aq) + Cu(s) → Cu2+(aq) + Ag(s)
  • 2Ag+(aq) + 2e- → 2Ag(s) (gaining electrons)
  • Cu(s) → Cu2+(aq) + 2e- (losing electrons)
  • 2Ag+(aq) + Cu(s) → Cu2+(aq) + 2Ag(s)