Electrochemistry

• Electrochemical reactions


  Reactions where electrons are transferred from one species to another
• Oxidation number


  A number assigned to keep track of what loses electrons and what gains them
• Assigning oxidation numbers


  1. Elements in elemental form have oxidation number 0

  2. Oxidation number of monatomic ion is same as charge

  3. Nonmetals tend to have negative oxidation numbers

  4. Sum of oxidation numbers in neutral compound is 0

  5. Sum of oxidation numbers in polyatomic ion is the charge on the ion

  
• Balancing oxidation-reduction equations using half-reaction method

  1. Assign oxidation numbers

  2. Write oxidation and reduction half-reactions

  3. Balance each half-reaction

  4. Multiply half-reactions to get same electrons gained/lost

  5. Add half-reactions, subtracting things on both sides

  6. Check equation is balanced for mass and charge

  
• Balancing oxidation-reduction equations in basic solution involves adding OH- to neutralize H+ and create water
• Voltaic cell


  A setup that uses a spontaneous oxidation-reduction reaction to do work by making electrons flow through an external device
• Voltaic cell


  • Oxidation occurs at the anode
  • Reduction occurs at the cathode
  • Salt bridge keeps charges balanced
• Potential difference


  The energy required to move a unit of electrical charge from one point to another against the electrostatic field
• Cell potential (emf or Ecell)


  The potential difference between the anode and cathode, the driving force for electrons to move in the external circuit
• Standard cell potential (Eocell)

  The cell potential under standard conditions of 25°C, 1M concentrations, 1 atm pressure
• Standard reduction potential (Eored)


  The potential associated with each electrode, the potential for reduction to occur at that electrode
• Cell potential is the difference between the standard reduction potentials of the cathode and anode reactions
• The standard hydrogen electrode (SHE) has a reduction potential of 0
• Standard state conditions


  o indicates standard state conditions
• Standard Reduction (Half Cell) Potential


  • We can tabulate the standard cell potential for all the possible cathode / anode combinations
  • It is not really needed to do so
• Cell potential


  • The difference between two electrode potentials
  • By convention the potential associated with each electrode is the potential for reduction to occur at that electrode
  • Standard electrode potentials are the standard reduction potentials denoted Eo

  red
• • Reduction potentials for many electrodes have been measured and tabulated
  • The standard Hydrogen electrode (SHE) has a potential of 0
• Standard Hydrogen Electrode


  • Their values are referenced to a standard hydrogen electrode (SHE)
  • By definition, the reduction potential for hydrogen is 0 V: 2 H+ (aq, 1M) + 2 e− → H2 (g, 1 atm)

  
• Voltaic cell based on half-reactions

  1. Determine Eo

  red for reduction of In3+ to In+

  2. Determine Eo

  red for reduction of Cu2+ to Cu
• Cell potential is an intensive property
• In any voltaic cell the reaction at the cathode has more positive value of Eo

  red than does the reaction at the anode
• Oxidizing and reducing agents


  The greater the difference between the two, the greater the voltage of the cell
• Strength of oxidizing and reducing agents

  • The more positive the Eo

  red value, the greater the tendency of the reactant to be reduced and therefore to oxidize other species

  • Zn(s) + Cu 2+(aq) Zn 2+(aq) + Cu(s), Eo

  red for Cu 2+ is 0.34 V, Eo

  red for Zn is -0.76 V
• Frequently used oxidizing agents

  • Halogens
  • O2
  • Oxianions whose central atoms have high oxidation states (MnO4

  –, Cr2O7

  1. , NO3

  • )
• Frequently used reducing agents

  • H2
  • Active metals (Alkali and Alkaline earth metals)
  • Zn and Iron (metals with negative Eo

  red values)

  
• Oxidizing and reducing agents

  • The strongest oxidizers have the most positive reduction potentials
  • The strongest reducers have the most negative reduction potentials
• Voltaic Cells use redox reactions that occur spontaneously
• Determine spontaneity of redox reactions using standard reduction potentials


  1. Identify oxidation and reduction half-reactions

  2. Find Eo

  red values

  3. Use formula Eo = Eo

  red (reduction) - Eo

  red (oxidation)
• Activity series


  • Decreasing ease of oxidation, from active metals to noble metals
  • Any metal on the list can be oxidized by the ions of elements below it

  
• Gibbs free energy


  • ΔG = ΔH + TΔS, measure of spontaneity of a process at constant temperature and pressure
  • Relationship between ΔG and cell potential E
• Concentration cells


  Cell potential is non-zero when concentrations are different, even if the same substance is at both electrodes
• Applications of oxidation-reduction reactions

  • Batteries
  • Alkaline batteries
  • Hydrogen fuel cells
  • Corrosion and corrosion prevention