A Determine the oxidation number for each atom in the following molecules H2S P2O5 S8 SCl2 Na2SO3 6. SO4-2 7. NaH Cr2O7-2 SnBr4 10. Ba(OH)2 For practice before class

12/05/06 Electrochemistry 19.9-19.13 p 941-955 A Intersection 14 12/05/06 Electrochemistry 19.9-19.13 p 941-955

December in Studio S M Tu W Th F A 12/5 Exam 3 12/6 Studio 12/8 Polymers; check out 12/11 Poster session, paper due 12/12 final IS 12/13 In-class assignment 12/17 Review session 7-9 pm 12/19 Final exam 8-10 am

Watershed Poster Session Monday, December 11 in USB 2165 Board (4 ft x 4ft), easel, pins Set up by 1:10 and 3:10 One person stationed at poster; others evaluate Rubric available Paper due same time

Last In-Class Assignment Wednesday, December 13th in studio Available on-line Read papers before coming to class; bring them with you. May make any notes you like on the papers Goal: to evaluate scientific method and data

Outline Ed’s demos Balancing Redox Reactions Electrochemistry Electrochemical cells and Standard Hydrogen Electrodes Nernst Quantifying current

A Ed’s Demos

Oxidation States of Vanadium: Reduction of V5+ to V+2 Reaction 1 Zn (s) + 2 VO3- (aq) + 8 H3O+ (aq) ↔ 2 VO2+ (aq) + Zn+2 (aq) + 12 H2O (l) Reaction 2 Zn (s) + 2 VO2+ (aq) + 8 H3O+ (aq) ↔ 2 V3+ (aq) + Zn+2 (aq) + 6 H2O (l) Reaction 3 Zn (s) + 2 V3+ (aq) ↔ 2 V2+ + Zn+2 (aq) V+5 (aq) → V+4 (aq) yellow to green V+4 (aq) → V+3 (aq) green to blue V+3 (aq) → V+2 (aq) blue to violet

Oxidation States of Manganese: Mn+7, Mn+6, Mn+4, and Mn+2 +7 (purple) to +2 (colorless) 2 MnO4- (aq) + H+ (aq) + 5 HSO3- (aq) ↔ 2 Mn+2 (aq) + 5 SO4-2 (aq) + 3 H2O(l) + 7 (purple) to +4 (brown) OH- + 2 MnO4- (aq) + 3 HSO3- (aq) ↔ 2 MnO2 (s) + 3 SO4-2 (aq) + 2 H2O(l) + 7 (purple) to + 6 (green) 2 MnO4- (aq) + 3 OH- + HSO3- (aq) ↔ 2 MnO4-2(aq) + SO4-2 (aq) + 2 H2O(l)

Thinking back…. What happened when Na(s) was added to water? Na(s) + H2O(l)  Na+ (aq) + H2(g) + OH-(aq) Determine the oxidation state of each reactant and product What was oxidized? What was reduced?

Balancing Redox Reactions M Balancing Redox Reactions When KMnO4 (potassium permanganate) is mixed with Na2C2O4 (sodium oxalate) under acidic conditions, Mn+2(aq) ions and CO2(g) form. The unbalanced chemical equation is: KMnO4(aq) + Na2C2O4(aq)  Mn+2(aq) + CO2(g) + K+(aq) + Na+(aq) K+ and Na+ are spectator ions, so we can ignore them at this point. MnO4- (aq) + C2O4-2(aq)  Mn+2(aq) + CO2(g)

Half-Reactions Reduction reaction Oxidation reaction M MnO4- (aq) + C2O4-2(aq)  Mn+2(aq) + CO2(g) Reduction reaction Oxidation reaction Reduction rxn Mn Oxidation C

Reduction reaction M MnO4-  Mn+2 Step 1: Balance all elements other than oxygen and hydrogen. Step 2: Balance the oxygens by adding water. Step 3: Balance the hydrogens using H+ Step 4: Balance the electrons Mn+7 on reactants side Mn+2 on products side Step 5: Check charge balance and elemental balance

M Oxidation reaction C2O4-2  CO2

Combine Half Reactions 5 e- + 8H+ + MnO4-  Mn+2 + 4 H2O C2O4-2  2 CO2 + 2e- 16 H+ + 2MnO4- +5C2O4-2 to 2Mn+2 + 8H2O + 10 CO2 We are assuming the reaction takes place under acidic conditions!

Balancing in Base 5 e- + 8H+ + MnO4-  Mn+2 + 4 H2O Change H+ to water by adding OH- to each side 5 e- + 8H+ + MnO4-  Mn+2 + 4 H2O C2O4-2  2 CO2 + 2e-

M NO2-(aq) + Cr2O7-2 → Cr+3(aq) + NO3-(aq) acidic soln

Electrochemical Cells

A Definitions Electrochemical cell: A combination of anode, cathode, and other materials arranged so that a product-favored redox reaction can cause a current to flow or an electric current can cause a reactant-favored redox reaction to occur Voltaic cell (battery): An electrochemical cell or group of cells in which a product-favored redox reaction is used to produce an electric current. Galvanic cell: A cell in which an irreversible chemical reaction produces electrical current Electrolytic cell: electrochemical reactions are produced by applying electrical energy

A Copper-Zinc battery – What Matters? Consider reduction potentials: Cu+2 + 2e- → Cu(s) 0.3419 V Zn+2 + 2e- → Zn(s) -0.7618 V Place Zn electrode in copper sulfate solution – What happens? Copper is plated on Zn electrode Cu+2 + 2e- → Cu(s) 0.3419 V Zn(s) → Zn+2 + 2e- 0.7618 V E > 0, spontaneous Cu+2 + Zn(s) → Zn+2 + Cu(s) 1.1 V Note, no need for electron to flow external to cell for reaction to occur!!

A Copper-Zinc battery – What Matters? Consider reduction potentials: Cu+2 + 2e- → Cu(s) 0.3419 V Zn+2 + 2e- → Zn(s) -0.7618 V Place Cu electrode in zinc sulfate solution – What happens? Zn doesn’t plate on copper electrode?! Cu(s) → Cu+2 + 2e- -0.3419 V Zn+2 + 2e- → Zn(s) -0.7618 V E < 0, not spontaneous Zn+2 + Cu(s) → Cu+2 + Zn(s) -1.1 V No reaction occurs !!

A Fig. 19-3, p.918

A                                                                                             What are the ½ reactions? What is the overall reaction? Identify the oxidation, reduction, anode, and cathode

SHE: Standard Hydrogen Electrode 2 H3O+(aq, 1.00 M) + 2e- <-> H2(g, 1 atm) + 2H2O(l) Eo = 0V Standard conditions: 1M, 1atm, 25oC Redox Reactions & Galvanic Cells In order to compare the potential difference when different metals are used for the anode and the cathode, each half reaction is compared to a standard half reaction.  This process is like choosing par in golf.  The par for a hole in golf is set at the standard.  Golfers scores are then measured relative to par. The freedom to set the cell potential difference equal to zero for this reaction is the same freedom we have to set any height equal zero wherever is convenient when we work with gravitational potential energy. We have this freedom because we are not interested in the absolute potential of a reaction, rather we are only interested in the magnitude of the change in potential. Fig. 19-7, p.922

Measuring Relative Potentials                                                                             Measuring Relative Potentials Table of Standard Reduction Potentials

Standard Reduction Potentials What is the standard potential of a Au+3/Au/Mg+2/Mg cell?

A The half-reaction with the more positive standard reduction potential occurs at the cathode as reduction. The half-reaction with the more negative standard reduction potential occurs at the anode as oxidation.