Voltaic Cells

What is Electrochem? The relationship between chemical reactions and electricity. Chemical energy converted to electrical energy Electrical energy converted to chemical energy

Electrochemical Cells Two types Cell produces electric current through a spontaneous reaction (galvanic aka voltaic) Cell consumes electric current to push a non- spontaneous reaction (electrolytic cells)

Components: A. Electrodes Anode Cathode B. Solutions C. Voltmeter D. Salt Bridge

Oxidation always occurs at the anode. Reduction always occurs at the cathode. Zn half reaction lies further to the right Why? Greater tendency to ionize than Cu. Makes Zn negatively charged (anode) relative to Cu electrode (cathode). Reaction driven by standard cell potentials.

Salt Bridges U-shaped tube with permeable stoppers. Contains strong electrolytes (KCl, KNO3, etc.) usually suspended in gel. Stops charge build-up in solutions for cells.

Standard Conditions Pressure = 1 atm for gaseous reactants. Temperature = assume 25oC Concentration= 1 M for each substance in solution

Cell Notation Write oxidation half reaction on left and reduction half reaction on right. Double vertical line (salt bridge) separates both half reactions Different phases are separated by a vertical line, same phases separated with a comma

Electromotive Force Emf = cell voltage = Eocell Difference in potentials that gives rise to an electric current Eocell = Volt (V) = J/C Coulomb = quantity of charge that passes a point in 1 sec when a current of 1 ampere flows.

Standard Cell Potentials Eocell = Eocathode + Eoanode Or Eocell = Eored(cathode) - Eored(anode) Eocell = positive for spontaneous reactions Eocell= negative for non-spontaneous reactions *All cell standards are reduction potentials. For oxidations, flip reaction and change the sign of potential.

Inert Electrodes Pt and graphite are typically electrodes for gas phase and liquid reactions. Ex. Standard Hydrogen Electrode (SHE) = uses Pt electrode. SHE consists of Pt electrode in contact with 1.00 M acid solution and H2 gas at 1 atm pressure. H2(g)  2H+(aq) + 2e- Eored = 0.00 V

Standard Electrode Potential The individual potential for each half reaction. SHE used as standard for calculating potentials of other half reactions. More negative charge = anode = more likely to undergo oxidation (think repelling e-s) More positive charge = cathode = more likely to undergo reduction (think attracting e-s)

Strength of oxidizing agents F2(g) + 2e-  2F-(aq)= most positive reduction potential Easiest to reduce…why? F2 = strongest oxidizing agent.

Strength of reducing agents Li+ +e-  Li(s) most negative Eored of -3.04 V Easiest to oxidize…why? Strongest reducing agent Wants to lose electrons, so reverse reaction occurs. Li(s)  Li+ +e- Eoox of +3.04 V

Practice Based on their reduction potentials, determine the best oxidizing agent and best reducing agent. Au3+ + 3e-  Au(s) 1.50 V Br2(l) + 2e-  2Br-(aq) 1.07 V Pb2+ + 2e-  Pb(s) -0.13 V Ni2+ + 2e-  Ni(s) -0.25 V Highest positive charge = Au = wants to gain electrons = best oxidizing agent.