Preliminary Concepts. Preliminary Concepts. Class 8.3 Oxidation/Reduction Reactions and Electrochemistry I. Friday, October 15 Chem 462 T.

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1 Class 8.3 Oxidation/Reduction Reactions and Electrochemistry I Friday, October 15 Chem 462 T. Hughbanks Preliminary Concepts Electrochemistry: the electrical generation of, or electrical exploitation of oxidation reduction reactions. Electrochemical reactions involve some means of producing or consuming electrons from an external source. The reacting system is an electrochemical cell and the electrical current enters or exits via electrodes. Preliminary Concepts Redox chemistry vs. electrochemistry: the former can occur independently of any electrochemical cell. Although we often use information from electrochemical experiments to understand redox processes, it is very often (usually?) the case that we are not interested in electrochemistry per se when discussing redox reactions.

2 More Preliminary Concepts Electrodes: Reduction occurs at one electrode (Cathode) Oxidation occurs at the other electrode (Anode) Electrical Conduction: movement of electrical charges from one place to another - usually through some medium. If the medium is a wire, it is metallic conduction; if the medium is an electrolyte solution, the conduction is carried out by ions. Electrochemical Cells Two kinds: Electrolytic: electrical current is used to drive otherwise nonspontaneous oxidationreduction (redox) reactions Galvanic (Voltaic): Spontaneous redox reactions are used to create electrical current (and do electrical work). Redox Reactions Oxidation-Reduction Reactions are ones in which electrons are transferred, eg.: CuSO 4 (aq) + Zn(s) ZnSO 4 (aq) + Cu(s) CuSO 4 dissolves in water to give a blue solution containing Cu 2+ and SO 2-4 ions. The reaction is: Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) This reaction is really an equilibrium, but K eq is huge: G rxn = G f,products - G f,reactants = kj/mol

3 Daniell Cell Zn Zn 2+ (1.0M) Cu 2+ (1.0M) Cu Separation in a Galvanic cell: Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) Cathode reaction: Cu 2+ (aq) + 2 e - Cu(s) Anode reaction: Zn(s) Zn 2+ (aq) + 2 e - Note: Despite what the picture seems to imply, electrons don t carry the current in solution! Galvanic (Voltaic) cells Zn Zn 2+ (1.0M) Cu 2+ (1.0M) Cu The Free Energy of a redox reaction ( G rxn ) can be harnessed by separating the oxidizing and reducing agents and forcing the electrons through a load. A Second Redox Reaction With a strong oxidant, Cu can donate electrons: 2 AgNO 3 (aq) + Cu(s) Cu(NO 3 ) 2 (aq) + 2 Ag(s) The net ionic reaction is: 2 Ag + (aq) + Cu(s) Cu 2+ (aq) + 2 Ag(s) Again, the reaction is really an equilibrium, but K eq is huge: G rxn = G f,products - G f,reactants = kj

4 Half Reactions Half-reactions: Ag + (aq) + e - Ag(s) Cathode reaction (reduction reaction) Cu(s) Cu 2+ (aq) + 2 e - Anode reaction (oxidation reaction) Again, the reactions can be separated in a Galvanic cell. Cu/Cu 2+ anode, Ag/Ag + cathode Cu Cu 2+ (1.0M) Ag + (1.0M) Ag 2 Ag + (aq) + Cu(s) Cu 2+ (aq) + 2 Ag(s) G rxn = G f,products - G f,reactants = kj Additivity of Cell Potentials 2 Ag + (aq) + Zn(s) Zn 2+ (aq) + 2 Ag(s) E cell = V V = 1.56 V (wired in series) G rxn = G f,products - G f,reactants = kj

5 Standard Hydrogen Electrode, SHE Pt H + (1.0M), H 2 (1 atm) Cu 2+ (1.0M) Cu H 2 (g) + Cu 2+ (aq) 2 H + (aq) + Cu(s) E cell = V All potentials can be referenced to one standard electrode. That electrode is the Standard Hydrogen Electrode shown here. Standard Reduction Potentials Reduction Half-reaction E (V) Ag + (aq) + e - Ag(s) 0.80 Cu 2+ (aq) + 2e - Cu(s) H + (aq) + 2e - H 2 (g) 0.00 (by defn.) Zn 2+ (aq) + 2e - Zn(s) The choice of SHE sets the zero of the scale, but all the measurable cell potentials don t depend on that choice. From Red. Potentials to Cell Potentials Reduction Half-reaction E (V) Cu 2+ (aq) + 2e - Cu(s) 0.34 Zn(s) Zn 2+ (aq) + 2e - [-0.76] Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) 1.10 The choice of SHE sets the zero of the scale, but all the measurable cell potentials don t depend on tis choice.

6 Using the Electromotive Series - Recipe for Evaluating Redox Rxns Spontaneity Find the appropriate half-reactions Write the half-reactions with the most positive (or least negative ) value first. Write the other half-rxn as an oxidation and write its oxidation potential (= reduction Potential) Balance the half-rxns. with respect to e transfer. (Don t multiply the potentials by the multiplicative constant used in balancing.) Add half-rxns and E s to get E cell, if > 0, it s spontaneous. Examples Can Ag(s) reduce Mg 2+ (aq) to metallic Mg (with formation of Ag + (aq))? Ag + (aq) + e - Ag(s) E = Mg 2+ (aq) + 2e - Mg(s) E = [ Ag + (aq) + e - Ag(s) ] E = Mg(s) Mg 2+ (aq) + 2e - E = Mg(s) + 2Ag + (aq) 2Ag(s) + Mg 2+ (aq) E cell = Examples Can MnO 4 oxidize Cl to Cl 2 to form Mn 2+ in acidic solution (all species 1.0 M)? E 2[ MnO H e Mn H 2 O ] [ Cl Cl e ] -(+1.36) 2MnO H Cl 2Mn H 2 O + 5Cl = E cell Note: H + (H 3 O + ) participates in the reaction. As we shall see quantitatively later, this means that the spontaneity of this reaction depends on ph. How do you think it affects matters?

7 More Definitions and Concepts Faraday: the amount of electricity in one mole of electrons: 1 Faraday = N A e = ( mol -1 )( C) 1 F = 96,485 Coulombs/mol Ampere: current flow equal to the passage of 1.0 C per second: 1 A = 1 C/s More... With the definitions above, we can quickly see that: If s amperes of current flow for t seconds, then st Coulombs of charge have passed through the circuit, and (st)/ 96,485 = # of moles of electrons that have passed through the circuit Key Relationship between G rxn and E G rxn is equal to the maximum amount of nonexpansion work that a reaction can do at const. T and P. In an electrical cell, this is the electrical work, w elec, that can be done: G rxn = w elec The electrical work to move a particle of charge q through a potential difference E is just qe. To move n moles of electrons (nn A electrons, each with charge -e) through a potential difference E, w elec = nn A (-ee) = -nfe G rxn = -nfe n is number of moles of electrons in the half-reactions of a cell reaction.

8 Some checks on units From electricity, we know Ohm s Law, V= IR J/ s ~ Power (Watts) ~ I 2 R = IV ~ Amps Volts J/ s = Amps Volts = C / s V J = C V Joules = Coulombs Volts G rxn = -nfe Check: Joules ~ mol Coulombs / mol Volts E cell and G rxn Very often, we refer to standard conditions for a cell reaction: G rxn = -nfe cell Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) E cell = 1.10 V All conditions are standard: T = 298 K, P = 1 atm, [Cu 2+ ] = [Zn 2+ ] = 1.0 M G rxn = -nfe cell = (-2)(96,485 C)(1.10 V) = kj From nonequilibrium to equilibrium Standard conditions, noneq.: G = G rxn < 0 E cell > 0 Final, equilibrium: G = 0 E cell = 0

9 Going to Equilibrium from Standard Conditions G = G rxn + RTlnQ (see Chap. 9) Starting at standard conditions, Q = 1, G = G rxn As we approach equil., Q K eq, G 0 we recover, G rxn = -RTlnK eq G rxn = -nfe cell Substitution gives the famous Nernst Eqn: E cell = E cell RT nf lnq as Q K eq, E cell 0 E cell = RT nf ln K eq Nernst Equation The Nernst Eqn., using base 10 logarithm: E cell = E cell RT nf log Q Often, 2.303RT is combined (T = 298 K) F E cell = E cell logq n Example 1 What is the initial potential, E cell? What will happen as the cell goes to equilibrium?