Lecture 244/1/05

Quiz 1) Balance the following redox equation: Ag(s) + NO 3 -  NO 2 (g) + Ag + (aq) 2) What is the oxidation number for Chlorine in ClO 4 - ?

Anode: Fe (s)  Fe 2+ (aq) + 2e - Cathode: 2e - + Cu 2+ (aq)  Cu (s) Net reaction: Fe(s) + Cu 2+ (aq)  Fe 2+ (aq) + Cu (s) (can solve and write like any other redox reaction) Electrochemical cell notation: Fe(s) | Fe 2+ (aq) ║ Cu 2+ (aq) | Cu (s)

Electromotive force (emf) or Cell potential or Cell voltage Measured by inserting a voltmeter into circuit produced by the difference in electrical potential between 2 electrodes Voltage determined by: identity of substances involved in redox concentration total moles of substances is not important for voltage

Electromotive force (emf) or Cell potential or Cell voltage E cell E° cell if at standard conditions 1 M of all the substances (1 atm if gases) 298 K Units are volts

If E° cell is positive, then reaction is product favored If E° cell is negative, then reaction is reactant favored E° cell = E° red + E° ox Reduction potential (this is the Cathode) Oxidation potential (this is the Anode)

Individual electrode potentials Can only measure potential differences, so need a reference reaction Standard Hydrogen Electrode (SHE) Hydrogen gas at a pressure of 1 atm bubbled over a platinum electrode immersed in 1 M aqueous acid Arbitrarily given E o red = 0 2 H 3 O + (aq, 1 M) + 2e -  H 2 (g, 1 atm) + 2 H 2 O

All other potentials determined from the SHE Zn (s)  Zn 2+ (aq, 1 M) + 2e - E° ox = ? 2 H 3 O + (aq, 1 M) + 2e -  H 2 (g, 1 atm) + 2 H 2 OE° red = 0 V__ Zn (s) + 2 H 3 O +  Zn 2+ + H H 2 O E° cell = V empirically determined with voltmeter Cathode: Arbitrarily assigned 0 Volts Anode E° oxidation = V

For any reaction: E° ox = - E° red Zn (s)  Zn 2+ (aq, 1 M) + 2e - E° ox = V OR Zn 2+ (aq, 1 M) + 2e -  Zn (s) E° red = V SO, if E° cell = E° red + E° ox Then E° cell = E° red (cathode) - E° red (anode) In practice, you will use the second equation

Standard Reduction Potentials 1. Each half-reaction is written as a reduction 2. Each half-reaction can occur in either direction 3. The more positive E° red, the more easily the substance on the left side of the half-reaction can be reduced 4. The more negative the E° red, the less likely the reaction will occur as reduction, and the more likely oxidation will occur 5. Under standard conditions,any species on the left side of a half-reaction will oxidize any species on the right that is farther down the table 6. Electrode potentials depend on the identity and concentration of the reactants and products, not the total quantity of each