This preview has intentionally blurred sections. Sign up to view the full version.
View Full DocumentThis preview has intentionally blurred sections. Sign up to view the full version.
View Full DocumentThis preview has intentionally blurred sections. Sign up to view the full version.
View Full Document
Unformatted text preview: Restricted: For students enrolled in Chem130/MCB100A, UC Berkeley, Fall 2006 ONLY Lecture 26 1 John Kuriyan: University of California, Berkeley Chem C130/MCB 100A, Fall 2006, Lecture 26 Review redox/potentials For such an electrochemical cell, under standard conditions, the voltage across the cell is the standard potential or voltage, E ° . The magnitude of the work done (electrical work) is charge x voltage. The charge moved per mole of reaction =  υ F where υ is the stoichiometry of the electrons and F is the Faraday constant (96500 C mol1 ). Hence the work done is  υ FE ° If there is no dissipation of heat, then the work done by the system is equal to the free energy change: Δ G ° =  υ FE ° (under standard conditions) Restricted: For students enrolled in Chem130/MCB100A, UC Berkeley, Fall 2006 ONLY Lecture 26 2 E ° values (reduction potentials) for several metals and important redox molecules are given in Tables. Reconsider the “lemon experiment” Restricted: For students enrolled in Chem130/MCB100A, UC Berkeley, Fall 2006 ONLY Lecture 26 3 Now we keep the negative pole of the voltmeter on the zinc electrode, but swaqp the copper electrode with an aluminum one. Δ G ° = RT ln K where K is the equilibrium constant for the reaction. Hence, we can calculate K if we know E ° . For the Zn/Cu battery, the reactions are Restricted: For students enrolled in Chem130/MCB100A, UC Berkeley, Fall 2006 ONLY Lecture 26 4 Zn → Zn 2+ + 2e E ° = + 0.762 V Cu 2+ + 2e → Cu E ° = + 0.34 V Cu 2+ + Zn → Zn 2+ + Cu E ° = + 1.10 V The equilibrium constant K = K = Zn 2 + ⎡ ⎣ ⎤ ⎦ Cu [ ] Cu 2 + ⎡ ⎣ ⎤ ⎦ Zn [ ] = Zn 2 + ⎡ ⎣ ⎤ ⎦ Cu 2 + ⎡ ⎣ ⎤ ⎦ (at equilibrium) The concentration of the pure metal is assumed to stay constant, at the standard state. Hence, knowing the value of E ° we can calculate the value of Zn 2 + ⎡ ⎣ ⎤ ⎦ Cu 2 + ⎡ ⎣ ⎤ ⎦ at equilibrium. Alternatively, if the system is not at equilibrium we can calculate the voltage generated by the system, using the Nernst equation. The free energy difference between the left and right sides of the cell is given by: Δ G = Δ G ° + RT ln Q divide both sides by υ F: Δ G υ F = Δ G υ F + RT υ F ln Q Restricted: For students enrolled in Chem130/MCB100A, UC Berkeley, Fall 2006 ONLY Lecture 26 5 ⇒ − E = − E + RT υ F ln Q ⇒ E = E − RT υ F ln Q (Nernst Eqation) where Q is the reaction quotient....
View
Full Document
 Fall '09
 Kuryian
 Electrochemistry

Click to edit the document details