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Electrochemistry. Redox reactions. Write the balanced equation for the reaction of permanganate and oxalate in acidic solution. MnO 4 - + C 2 O 4 2-  Mn 2+ + CO 2 1. Find oxidation numbers. Permanganate & oxalate. Separate into half reactions

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Redox reactions

Redox reactions

  • Write the balanced equation for the reaction of permanganate and oxalate in acidic solution.

    • MnO4- + C2O42- Mn2+ + CO2

  • 1. Find oxidation numbers

Permanganate oxalate

Permanganate & oxalate

  • Separate into half reactions

  • Balance for mass, balance O with H2O, balance H with H+, balance charge w/ electrons

Finishing the problem

Finishing the problem

  • Balancing the electrons

  • Combine and cancel

Redox titration

Redox titration

  • Use the ratio from the balanced redox reaction to determine concentrations.

  • How will we know when the reaction is complete?

    • We use an indicator

    • In our earlier redox reaction, permanganate is a “self indicator”.

Activity series

Activity series

  • Some substances are more easily oxidized than others.

  • Listing these in order of their ability to oxidize gives us an activity series.

  • A substance will oxidize if its above the item reducing it on the table.

    • Cu will be oxidized by Ag+ because its higher on the activity series. Ag cannot be oxidized by Cu2+.

Galvanic cell

Galvanic Cell

  • Using a spontaneous redox reaction to generate electricity.

  • Anode = oxidation (AnOx)

  • Reduction at the cathode (RedCat)

  • Electrons flow From Anode To Cathode (FatCat)

Salt bridge

Salt Bridge

  • Allows ions to travel between solutions, maintaining charge balance.

  • Can be a solution, a paste or a porous membrane

Galvanic cell notation

Galvanic cell notation

  • A shorthand way of writing the structure of a galvanic cell is

  • Anode | Anode ion || Cathode Ion | Cathode

  • | denotes a phase barrier (solid metal and aqueous solution)

  • || denotes the salt bridge

  • Mg(s) | Mg2+(aq) || Al3+(aq) | Al(s) is the cell notation of an aluminum/magnesium cell

Cell potential

Cell potential

  • The Standard Reduction Potential table measures the voltage produced when that half reaction is reduced by the standard hydrogen electrode

    • A negative voltage means that electricity must be applied to cause the reaction (more later)

  • All concentrations are assumed 1M, pressures are 1 atm at 298K.

  • How does this compare to the activity series?

Cell potential of a galvanic cell

Cell potential of a galvanic cell

  • Ecell = Ered + Eox

  • When you flip the oxidation half reaction, change the sign of the voltage!!!

  • What is the cell potential of a copper/zinc cell?

Still more gibbs and some equilibrium

Still more gibbs (and some equilibrium)

  • ΔG° = -nFE°

    • n is the number of moles of electrons transferred (from balanced equation)

    • F is the Faraday constant (~96,500 coulomb/mol)

    • E° is the standard cell potential



  • Calculate ΔG° and K from the following reaction:

  • 4Ag(s) + O2(g) + 4H+(aq)  4Ag+(aq) + 2H2O(l)

Cell potential at nonstandard conditions

cell potential at nonstandard conditions

  • Gibbs at Nonstandard Conditions

    • ΔG = ΔG° + RT lnQ

  • Substituting –nFE into the equation gives us the Nernst Equation:

    At 298K, we can rewrite it as

Practical use of nonstandard conditions

Practical use of nonstandard conditions

  • Concentration cell

    • Same anode and cathode

    • Same electrolyte

  • E° of the cell is 0.

  • Due to the huge differences in concentration, Q will not = 1. This gives a small voltage.

Concentration cell

concentration cell

  • What voltage will be generated by the cell on the right?



Hydrogen fuel cells

hydrogen Fuel Cells

  • Hydrogen is oxidized to form H+ and e-.

  • The electron passes through a wire and performs work.

  • H+ travels through the membrane, combines with O2 and the electron to form H2O.



  • Most metals allow free flow of electrons.

  • In the presence of H+, Fe will oxidize to Fe2+.

  • Water acts as both a source of H+ and the electrolyte for Fe2+ ions to flow.

  • Further oxidation of Fe2+ and combination with oxygen form Fe2O3.

Preventing corrosion

Preventing corrosion

  • Galvanization

    • Coating a corrodible metal will prevent corrosion.

  • Cathodic Protection

    • Attaching a metal that is easier to oxidize will prevent corrosion.

      • Ships attach large pieces of magnesium to their hull. The magnesium will oxidize in place of the iron.



  • Adding electricity to a nonspontaneous reaction will cause it to occur.

  • We can generate sodium metal and chlorine gas from a molten NaCl sample.

Predicting electrolytic half reactions

predicting electrolytic half reactions

  • Use your reduction potential table. The combination that produces the most positive voltage will occur.

  • Beware of aqueous solutions

    • Water can be oxidized to form O2, or reduced to form H2. You must keep these in mind when figuring out what reactions occur.

Molten vs aqueous

molten vs aqueous

  • Molten NaI will form metallic sodium and iodine. Will NaI(aq) do the same?

  • There are four possible reactions:

    • Oxidation

      • 2I- I2 + 2e- E = -0.54V

      • 2H2O  O2 + 4H+ + 4e- E = -1.23V

    • Reduction

      • Na+ + e-  Na E = -2.71V

      • 2H2O + 2e-  H2 + 2OH- E = -0.83V

  • Let’s prove it!



  • If we have a metallic anode and an electrolyte solution of that metal, we can coat an object (serving as the cathode) with that metal.

Electroplating stoichiometry

electroplating stoichiometry

  • We know that:

    • 1 mol e- = 96,500 coulombs (Faraday’s Constant)

    • 1 Amp = 1 coulomb/sec

  • Electroplating questions are simple factor-label!



  • Calculate the number of grams of aluminum produced in 1.00 hour by the electrolysis of molten AlCl3 at a current of 10.0 A.

  • Al3+ + 3e- Al

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