Electrochemistry Chapter 20 Brown, LeMay, and Bursten

Definition  The study of the relationships between electricity and chemistry  Review redox reactions  Review balancing redox reactions in acid and base

Voltaic Cell (also called Galvanic Cell)  Device in which the transfer of electrons takes place through an external pathway.  Electrons used to do work

Summary of Cell  Each side is a half-cell  Electrons flow from oxidation side to reduction side – determine which is which  Salt bridge allows ions to move to each terminal so that a charge build up does not occur.  Assignment of sign is this: Negative terminal = oxidation (anode) Positive terminal = reduction (cathode)  Salt bridge allows ions to move to each terminal so that a charge build up does not occur. This completes the circuit.

Cell EMF  Flow is spontaneous  Caused by potential difference of two half cells. (Higher PE in anode.)  Measured in volts (V)  1 volt = 1 Joule/coulomb  This is the electromotive force EMF (force causing motion of electrons through the circuit.

E cell  Also called the cell potential, or E cell  Determined by reactant types, concentrations, temperature  Under standard conditions, this is E° cell  25° C, 1 M or 1 atm pressure  This is 1.10 V for Zn-Cu  Shorthand: Zn/Zn 2+ //Cu 2+ /Cu

Reduction Potentials  Compare all half cells to a standard (like sea level)  2H + + 2e - → H 2 (g) = 0 volts (SHE)  The greater the E°red, the greater the driving force for reduction (better the oxidizing agent)  In a sense, this causes the reaction at the anode to run in reverse, as an oxidation.  Use this equation:  E°cell = E°red (cathode) - E°red (anode)

Trends

Spontaneity  Positive E value indicates that the process is spontaneous as written.  Activity series of Metals – listed as oxidation reactions  Reduction potentials in reverse  Example, Ag is below Ni because solid Ni can replace Ag in a compound. Actually, Ni is losing electrons and thus being oxidized by Ag +. Ag is listed very high as a reduction potential.

Relationship to ΔG  ΔG = -nFE n = number of electrons transferred F = Faraday constant = 96,500 C/mol or 96,500 J/V-mol  Why negative? Spontaneous reactions have +E and – ΔG.  Volts cancel, units for ΔG are J/mol  Standard conditions: ΔG° = -nFE°

Nernst Equation  Nonstandard conditions – during the life of the cell this is most common  Derivation  E = E ° - (RT/nF)lnQ  Consider Zn(s) + Cu 2+ → Zn 2+ + Cu(s)  What is Q?  What is E when the ions are both 1M?  What happens as Cu 2+ decreases?

Concentration Cells  Same electrodes and solutions, different molarities.  How will this generate a voltage? Look at Nernst Equation. E = E ° - (RT/nF)lnQ  When will it stop?  Basis for a pH meter and regulation of heartbeat in mammals

EMF and equilibrium  When cell continues to discharge, E eventually reaches 0. At this point, because ΔG = -nFE, it follows that ΔG = 0.  Equilibrium!  Therefore, Q = Keq  Derivation  logKeq = nE°/0.0592

Batteries  Portable, self-contained electrochemical power source  Batteries in series, voltage is added.

Things to consider  Size (car vs. heart)  Amount of substances before it reaches equilibrium  Toxicity (car vs. heart)  A lot a voltage or a little (car vs. heart)  Example – alkaline camera battery  Dry – no water

Fuel Cells  Not exactly a battery, because it is open to the atmosphere  How does the combustion of fuel generate electricity? – heats water to steam which mechanically powers a turbine that drives a generator – 40% efficient  Voltaic cells are much more efficient  l8.swf l8.swf

Corrosion  Undesirable spontaneous redox reactions  Thin coating can protect some metals (like aluminum) – forms a hydrated oxide)  Iron - $$$$$

Protection  Higher pH  Paint surface  Galvanize (zinc coating) – why?  Zinc is a better anode  Called cathodic protection – sacrificial metal

More dramatic

Electrolysis  Cells that use a battery or outside power source to drive an electrochemical reaction in reverse  Example NaCl → Na + + Cl -  Reduction at the cathode, oxidation at the anode  Voltage source pumps electrons to cathode.

Diagram

Solutions  High temperatures necessary for previous electrolysis (ionic solids have high MP)  Easier for solutions, but water must be considered  Example: NaF  Possible reductions are: Na + + e - → Na(s) (Ered = V) 2H 2 O + 2 e - → H 2 (g) + 2 OH- (Ered = -.83 V)  Far easier to reduce water!  continue

Continued  Look at possible oxidations: 2F- → F 2 (g) (Ered = 2.87 volts) 2H 2 O → O 2 (g) + 4H + + 4e - (Ered = 1.23 volts) Far easier to oxidize water, or even OH-!  So for NaF, neither electrode would produce anything useful, and doesn’t by experiment  With NaCL, neither electrode is favored over water. However, the oxidation of Cl- is kinetically favored, and thus occurs upon experimentation!  Use Ered values of two products to find Ecell (minimum amount of energy that must be provided to force cell to work)

Active electrodes  If electrode is not inert, it can be coated with a thin layer of the metal being reduced, if its reduction potential is greater than that of water.  This is called electroplating  Ecell = 0, so a small voltage is needed to push the reaction.

Quantitative relationship