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Discover the fundamentals of electrochemistry, including redox reactions, cell structures, EMF, reduction potentials, Nernst equation, and applications in batteries, fuel cells, and corrosion prevention.
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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.
