Chapter 17 Electrochemistry Redox review (4.9) 17.1-17.2 17.4-17.5 17.6-17.7 Review Oxidation-Reduction Involves transfer of electrons from reducing agent to oxidizing agent Oxidation= loss of e (increase in oxid #) Reduction= gain of e (decrease in oxid#) GER and LEO REVIEW 1. 2. 3. 4. 5. atom in element = 0 monatomic ion = charge fluorine = -1 oxygen = -2 hydrogen = +1 6. sum of oxid. # in compound = 0 7. sum of oxid. # in polyatomic ion = charge on ion The Half-Reaction Method (Acidic Solution) Copyright © Houghton Mifflin Company. All rights reserved. 4–4 Review- Balancing OxidationReduction Reactions 1. Separate in ½ reactions 2. Intermediate steps a. balance all elements other than H and O b. balance O with H2O c. balance H with H+ d. balance charge with (e-) 3. Multiply ½ rxn. so that the number of electrons is same 4. Add ½ rxns. Capture the Energy MnO4- + 5Fe2+ Mn2+ 5Fe3+ MnO4- and Fe2+ will react directly in solution. Electrons will be transferred and energy will be released as heat. No useful work will result. Capture the Energy! Zn + Cu2+ - Zn2+ + Cu Separate ½ reactions Connect metals w/ wire electrons) (to transfer Connect soln w/ bridge (keeps solns separate but allows ions to move) Converts Chemical Energy to Electrical Energy!!- A Battery!! Galvanic Cell Capture the energy You have separated the oxidizing agent from the reducing agent Requires electron transfer through wire Attach a motor, light bulb, bell etcthe current produced in the wire by eflow provides work!! Figure 17.6 A Galvanic Cell involving the Half-Reactions Copyright © Houghton Mifflin Company. All rights reserved. 17–10 Cell potential is….. The pressure of a Galvanic cell to “push” the e- “driving force” Electromotive Force, emf Symbol E Units: Joule/coulomb (=1Volt, V) Coulomb = unit of charge Specifies # of e- E cell =E anode +E cathode (oxidation) pushing e- (reduction) pulling e- (black wire) (red wire) A spontaneous rxn in a Galvanic cell must be positive. E>0 E 1/2 reactions P. 796 table Standard Reduction potentials 1M solutions 1atm gases 25 C Hydrogen ½ rxn = 0.00V Table 17.1 Standard Reduction Potentials at 25°C (298K) for Many Common Half-Reactions Copyright © Houghton Mifflin Company. All rights reserved. 17–14 Helpful Info Need balanced oxidation-reduction rxns from the reduction potentials. One reduction ½ rxn must be reversed. * The ½ rxn with largest positive potential will run as written (reduction). The other ½ rxn will run in reverse (oxidation). Reversing Direction Changes Sign of E Because: E oxidation = -E Then: E cell = E reduction cathode Examples: –E anode Standard Reduction Potential Math Rules # of e- lost must equal # e- gained ½ rxns must be multiplied by integers to balance equations Value of E is not changed when ½ rxn multiplied by an integer. Potential is NOT multiplied by integer. Example…. Line Notation Anode listed on left Cathode listed on right Mg(s) l Mg2+ ll Al3+l Al(s) Anode Mg0(s) - Mg2+ Cathode Al3+ - Al0(s) Cell Potential & Free Energy A galvanic cell will run in the direction that gives a positive value for E +E corresponds to - G +E and - G indicates a spontaneous reaction. G = -n FE G = nFE n = # of e(exchanged in overall rxn) F = 96,485(c/mol e-) (Faraday’s constant) Examples: Effects of Concentration on E So far the cells have been under standard conditions…. Le Chatelier’s principle applies if not std. conditions.. Determine if E cell > or < E cell ?? To summarize: If E cell not at standard conditions: [Reactant] > 1mol/L E cell > E *cell [Product] < 1mol/L E cell> E *cell Reverse is also true Concentration cell Same components in cells, but different concentrations. Equilibrium wants these concentrations to be Equal. Examples: Nernst Equation Establishes relationship b/t cell potential and concentration of cell components. For cells not at 1M Concentration: E = E * - RT/nF ln (Q) E * is std cell potential RT/nF ln (Q) is correction factor Common form: E = E * - RT/nF ln (Q) Commonly written : E = E * - 0.0591/n log (Q) Examples: A Battery @ Equilibrium At Equilibrium: Ecell = 0 (completely discharged) Q=K and delta G = 0 Using the Nernst Equation: @Equilibrium: 0 =E * - 0.0591/n log(K) Or log K = nE */0.0591 Corrosion Process of returning metals to their natural state. Metals oxidize readily resulting in corrosion. Metal ½ rxn is reversed for oxidation. Combined with Oxygen ½ rxn. to give (+) Ecell Electrolysis Involves forcing current through a cell to produce a chemical change resulting in (-) cell potential. Example: Figure 17.19 a-b (a) A Standard Galvanic Cell Based on the Spontaneous Reaction Zn + Cu2+ - Zn2+ + Cu (b) A Standard Electrolytic Cell. A Power Source Forces the Opposite Reaction Cu + Zn2+ - Cu2+ + Zn. Copyright © Houghton Mifflin Company. All rights reserved. 17–29