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Electrochemistry

Electrochemistry. Electrochemistry = the interchange of chemical and electrical energy = used constantly in batteries, chemical instruments, etc… Galvanic Cells Definitions

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Electrochemistry

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  1. Electrochemistry Electrochemistry = the interchange of chemical and electrical energy = used constantly in batteries, chemical instruments, etc… Galvanic Cells • Definitions • Redox Reaction = oxidation/reduction reaction = chemical reaction in which electrons are transferred from a reducing agent (which gets oxidized) to an oxidizing agent (which gets reduced) • Oxidation = loss of electron(s) to become more positively charged • Reduction = gain of electron(s) to become more negatively charged • Using Redox Reactions to generate electric current (moving electrons) • Zno + Cu2+ Cuo + Zn2+ • Zno is oxidized and Cu2+ is reduced • Half Reaction = oxidation or reduction process only Reduction: Cu2+ + 2e- Cuo Oxidation: Zno Zn2+ + 2e-) Sum = Redox Rxn

  2. 2) In solution: • Zno and Cu2+ collide and electrons are transferred • No work can be obtained; only heat is generated

  3. 3) In separate compartments, electrons must go through a wire = Galvanic Cell a) Generates a current = moving electrons from Zno side to Cu2+ side b) Current can produce work in a motor or light up a light bulb c) Salt Bridge = allows ion flow without mixing solutions (Jello-like matrix) d) Chemical reactions occur at Electrodes = conducting solid dipped into solution • Anode = electrode where oxidation occurs (production of e-) • Cathode = electrode where reduction occurs (using up e-)

  4. C. Cell Potential • Think of the Galvanic Cell as an oxidizing agent “pulling” electrons off of the reducing agent. The “pull” = Cell Potential • ecell = Cell Potential = Electromotive Force = emf • Units for ecell = Volt = V 1 V = 1 Joule/1 Coulomb • Voltmeter = instrument drawing current through a known resistance to find V Potentiometer = voltmeter that doesn’t effect V by measuring it • Standard Hydrogen Electrode: must have a standard to compare emf to Cathode = Pt electrode in 1 M H+ and 1 atm of H2(g) Half Reaction: 2H+ + 2e- H2(g) e1/2 = 0

  5. 4) Standard Reduction Potentials can be found in your text appendices • Always given as a reduction process • All solutes are 1M, gases = 1 atm 5) Combining Half Reactions to find Cell Potentials • Reverse one of the half reactions to an oxidation; this reverses the sign of e1/2 • Don’t need to multiply for coefficients = Intensive Property (color, flavor) • Example: 2Fe3+(aq) + Cuo 2Fe2+(aq) + Cu2+(aq) • Fe3+ + e- Fe2+ e1/2 = +0.77 V • Cu2+ + 2e- Cuoe1/2 = +0.34 V • Reverse of (ii) added to (i) = -0.34 V + +0.77 V = +0.43 V = e1/2

  6. Direction of electron flow in a cell • Cell always runs in a direction to produce a positive ecell • Fe2+ + 2e- Feoe1/2 = -0.44 V MnO4- + 5e- + 8H+ Mn2+ + 4H2O e1/2 = +1.51 V • We put the cell together to get a positive potential: DG = -nFecell • 5(Feo Fe2+ + 2e-)e1/2 = +0.44 V • 2(MnO4- + 5e- + 8H+ Mn2+ + 4H2O) e1/2 = +1.51 V 16H+(aq) + 2MnO4-(aq) + 5Feo(s) 2Mn2+(aq) + 5Fe2+(aq) + 8H2O(l) ecell = 1.95V

  7. V • Notes on the Experimental Procedure • Do all of Part I: Direct Redox Reactions • Part II Indirect Spontaneous Redox Reactions: Only do procedures 1-16 • Skip 17-20 of Part II and all of Part III: Indirect Non-Spontaneous Reactions • Make sure to clean metal electrodes with sandpaper to get best results • Potentials may not be identical to predicted, but the relative sizes will be • Use Nickel in place of Tin (Sn); it works better • Don’t throw away any metal pieces; clean them, dry them, put them back • Use the same piece of filter paper (NaNO3 soaked) for all galvanic cells • Place solutions in wells on plate, so that salt bridge can reach all needed

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