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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).
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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.
Ecell • Also called the cell potential, or Ecell • 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/Zn2+//Cu2+/Cu
Reduction Potentials • Compare all half cells to a standard (like sea level) • 2H+ + 2e-→ H2(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)
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) + Cu2+ → Zn2+ + Cu(s) • What is Q? • What is E when the ions are both 1M? • What happens as Cu2+ 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 • http://www.fueleconomy.gov/feg/fuelcell8.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
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.
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 = -2.71 V) • 2H2O + 2 e- → H2(g) + 2 OH- (Ered = -.83 V) • Far easier to reduce water! • continue
Continued • Look at possible oxidations: • 2F- → F2(g) (Ered = 2.87 volts) • 2H2O → O2(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.
