Download Chapter 11 Electrochemistry

Warren Mar
12/13/14
Chapter 11 Electrochemistry
11.1 Galvanic Cells
 Oxidation – loosing electrons ( increase in oxidation number) – occurs at anode.
 Reduction – gaining electrons ( decrease in oxidation number) – occurs at cathode.
 Galvanic (Voltaic) cell – chemical energy is changed to electrical energy.
 Salt bridge or porous disk needed to balances charges in a galvanic cell.
A. Cell Potential
 Electromotive force (cell potential) is measured in volts = joules of work per
coulomb of charge transferred. 1V = 1J/C.
 Voltmeter – measures voltage, but loses potential to heating wire.
 Potentiometer has no current flow, so maximum cell potential is measured.
11.2 Standard Reduction Potentials
 Standard hydrogen electrode – Pt electrode surrounded by H2 at 1atm. And in contract
with 1M H+ ions – arbitrarily assigned value of exactly 0V.
 Sign of is reversed E° (standard reduction potentials) if half-reaction is reversed.
 When half-reaction is multiplied E° does not change (intensive property – not
dependent on how many times a reaction occurred.)
A. Items Needed for a Description of a Galvanic Cell
 The cell potential (always positive for galvanic cell) and the balanced cell reaction
 The direction of electron flow, obtained by inspecti0on the half-reactions and using
the directions that give a positive E°cell.
 Designation of the anode and the cathode.
 The nature of each electrode and the ions present in each compartment. A
chemically inert conductor is required if none of the substances participating in the
half0 r3eaction is a conduction solid.
11.3 Cell Potential, Electrical Work, and Free Energy
 Potential difference (V) = emf = work (J)/ charge (C). E = -w/q, where w is work and q
is charge from the point of view of the system. wmax = -qEmax.
 Charge on 1 mole of electrons (faraday) = 96,485 coulombs of charge per mole of
electrons. q = nF.
 G = wmax = -nFEmax = -qEmax. G° = -nFE°. For a spontaneous reaction G is
negative, so it works out.
11.4 Dependence of the Cell Potential on Concentration
A. The Nernst Equation
 G = G° + RTln(Q). E = E° - RT/nF*ln(Q).
 Q = K and Ecell = 0. This is a dead battery at equilibrium after reaction has
completed.
B. Ion-Selective Electrodes
 Electrodes which are sensitive to concentrations of a particular ion.
 Glass Electrodes can change potential depending on concentration of ion.
C. Calculation of Equilibrium constants for Redox Reactions
 Assume Q = K and Ecell = 0. Solve for K.
D. Concentration Cells
 Same components different moralities.
 Plug in Q to solve for V.
11.5 Batteries
 Sources of direct current (DC) and can be connected in series to add to the total battery
potential.
A. Lead Storage Battery
 Lead as anode. Lead dioxide as cathode. Sulfuric Acid (H2SO4) as electrolyte.
 Each cell produces about 2 volts x6 = 12 volts.
 Jump starting can be dangerous due to electrolysis of water in dead battery which
produces H2. Disconnecting jumper cables can produce an arc unless you ground
cable to part of the engine away from battery. New batteries are sealed, because
they have electrodes, which inhibit electrolysis of water.
B. Dry Cell Battery
 Most contain zinc. Alkaline batteries last longer than acidic batteries, because Zn
corrodes less rapidly in basic conditions.
 NiCd can be recharge indefinitely like Pb batteries.
C. Fuel Cells
 Galvanic cell in which the reactants are continuously supplied.
11.6 Corrosion
 Sometimes an oxidation coating prevents metal from further corroding like Al.
A. Corrosion of Iron
 Moisture forms a salt bridge, that is why car last longer in the Southwest.
B. Prevention of Corrosion
 Protection by application of a coating, paint or metal plating. Chromium and Tin
often used to plate steel.
 Galvanizing, Zn used to coat steel.
 Alloying is another way to prevent corrosion, stainless steel for example.
 Surface alloy, new technology uses plasma directed onto surface of the metal.
 Cathodic protection – an active metal (Mg) is connected by wire to a pipe. Mg is a
better reducing agent, so it reacts instead of the pipe. Mg anode needs to replaced
over time.
11.7 Electrolysis
 Electrolytic cell – uses electrical energy to produce chemical change that would not
otherwise occur spontaneously.
 Ampere (amp) = 1 coulomb of charge per second.
A. Electrolysis of Water
 Pure water contains so few ions that only a small amount of current can flow.
 If soluble salts are introduced bubbles of H2 and O2 form.
B. Electrolysis of Mixtures of Ions

Selectively plating out ions. The greater the E, the more likely the reaction will
proceed in that direction making it easier to plate out.
 Overvoltage is voltage that is much higher than expected voltage needed to plate
out ion, which is due to difficulties in transferring electrons from the species in the
solution to the atoms on the electrode across the electrode solution interface.
11.8 Commercial Electrolytic Processes
A. Production of Aluminum
 Hall-Heroult process – uses molten cryolite (Na3AlF6) as the solvent for the
aluminum oxide. H2O reduces more readily than Al. Chemical process is not
completely understood.
B. Electrorefining of Metals
 Cathode is ultrapure metal and anode is impure. Pure metal is deposited at the
cathode.
C. Metal Plating
 Cathode is plated. For example chrome bumpers and “tin” cans, which are actually
steel.
D. Electrolysis of Sodium Chloride
 Downs cell used to produce NaOH and Cl2. Mercury cell electrolyzes aqueous
NaCl, brine.