Electrochemistry

• Oxidation


  A species (an atom or ion) loses electrons in a reaction
• Reduction


  A different atom or ion in the reaction gains the electrons lost
• Redox reaction


  An oxidation-reduction reaction
• Balancing redox reactions


  1. Balance any atom other than O or H

  2. Balance O by adding water

  3. Balance H by adding H+

  4. Balance charge by adding electrons
• Oxidation number


  • Elements have an oxidation number = zero

  2. Monoatomic ions equals the ion's charge

  3. Fluoride is -1

  4. Oxygen in a compound or ion is -2 except in peroxide (O22-) where each O is -1 and if O is bonded to F (OF2) then O is +2

  5. Hydrogen is +1 or if with a metal is -1 (H-1 is called hydride)

  6. Oxidation numbers in compounds must equal zero

  7. Oxidation numbers in a polyatomic ion must equal the charge of the ion

  8. Assign the most electronegative element a negative oxidation number
• Steps in balancing redox reactions


  Split the reaction into 2 half reactions

  2. Starting with one of the half reactions, balance all atoms other than O and H using coefficients

  3. Balance O by adding water to either side of equation as needed

  4. Balance H by adding H+ ion to either side of equation as needed

  5. Add up all charges on each side of the half reaction. Add as many electrons as needed to one side so that charges of reactants now equal charges on products side

  6. Go to the other half reaction and repeat steps #2-5

  7. Look at the electrons in both half reactions. The number of electrons lost much equal the number of electrons gained. If the number of electrons in each half reaction is not equal then multiply the entire half reaction by an integer that will make the electrons balanced

  8. Add the half reactions back together and the electrons should cancel since you should have the same coefficient in front of the electrons lost (on product side) and the electrons gained (on reactant side)

  9. Check to see that atoms and charges balance
• Balancing Redox Example #1


  • MnO4- + Fe2+ → Mn2+ + Fe3+
• Balancing Redox Example #2

  • Al3+(aq) + Mg (s) → Al (s) + Mg2+(aq)
• Electrochemistry is the interchange of chemical and electrical energy. The key to electrochemistry is a redox reaction.
• Redox reactions produce a flow of electrons that can be used to do work like light a flashlight bulb or provide energy for your cell phone to function.
• A battery or galvanic cell is a type of electrochemical cell.
• Oxidation Involves Loss of electrons; Reduction Involves Gain of electrons (OIL RIG) OR Losing Electrons is Oxidation; Gaining Electrons is Reduction (LEO the lion says GER)
• The law of conservation of mass is obeyed in redox reactions. The numbers of atoms must be balanced on both sides of the reaction, and the charges must also be balanced.
• Balancing Redox Reaction


  1. Split the reaction into half-reactions

  2. Balance atoms/ions

  3. Balance electrons

  4. Multiply half-reactions to equalise electrons

  5. Add half-reactions
• Galvanic Cell


  A battery, an electrochemical cell where a spontaneous redox reaction occurs to produce electrical energy
• Components of a Galvanic Cell


  • Anode (where oxidation occurs, mass decreases)
  • Cathode (where reduction occurs, mass increases)
  • Voltmeter (measures cell potential)
  • Salt bridge (maintains electrical neutrality)
• How a Galvanic Cell Operates


  1. All reactants and products present before reaction

  2. Electrodes made of metals or inert substances

  3. Determine anode and cathode from reduction potentials

  4. Electrons flow from anode to cathode

  5. Salt bridge allows flow of ions to maintain charge balance
• Standard Reduction Potentials in Aqueous Solutions at 25°C
• Determining anode and cathode

  More positive reduction potential is the cathode, the other is the anode
• Cell Potential (E°cell)


  Difference between reduction potentials of cathode and anode
• E°cell = E°cathode - E°anode
• Cell with higher potential represents the cathode, the smaller one is the anode
• Applications of Cell Potential


  • Determine direction of electron flow
  • Predict feasibility of reactions
  • Determine strength of reducing/oxidizing agents
  • Predict standard cell potential
• Gibb's Free Energy and Electrochemistry


  • Work accomplished = Potential difference (V) x Charge transferred (C)
  • 1 Joule of work = 1 Volt x 1 Coulomb
• If work flows out, it is assigned a negative sign
• Anode


  Larger electrode where oxidation occurs
• Application of cell potential


  • To determine the direction of electron flow
  • To predict the feasibility of reactions
  • To determine the strength of reducing agent and oxidizing agents
  • To predict the standard cell potential, E⁰cell
• Emf (V)


  The work that can be accomplished when electrons are transferred through a wire depends on the potential difference (in volts) between 2 points in the circuit
• Joule (J)


  One Joule of work is produced when one Coulomb of charge is transferred between two points in the circuit that differ by a potential of one volt
• Coulomb (C)


  A quantity of electrons or charge passing through the circuit
• If work flows OUT


  It is assigned a minus sign so that when a galvanic cell produces electricity the cell potential is positive and can be used to do work so work has a negative sign
• Gibb's Free Energy (G)


  Discussed in great detail in Thermodynamics. Its sign is very important, for a reaction to occur on its own or be thermodynamically favorable the sign of ∆G⁰ must be negative
• Batteries


  Chemical reactions that occur when components are in contact with one other so are thermodynamically favorable (∆G⁰ -)
• n


  The number of moles of electrons transferred (balanced number from the half-reactions that cancels out)
• F


  Faraday's constant (the charge of one mole of electrons) = 96500 C/mole of electrons
• E⁰


  The standard cell potential from E⁰ = E⁰reduction - E⁰oxidation
• The reaction is thermodynamically favorable since ∆G⁰ is negative
• Equilibrium constant (K)


  Describes how the reaction proceeds and is simply a ratio of concentrations. If a reaction like a galvanic cell is going forward, it produces a lot of products and uses the reactants so that the ratio of K > 1
• Since K>>1, the reaction is definitely going forward
• Characteristics of galvanic cells


  • E⁰ is positive
  • ∆G⁰ is negative
  • K > 1