Transcript CH 20: Electrochemistry

Electrochemistry
Chapter 20
Brown, LeMay, and Bursten
Definition
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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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Device in
which the
transfer of
electrons takes
place through
an external
pathway.
Electrons used
to do work
Summary of Cell
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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:
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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
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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
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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
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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:
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E°cell = E°red (cathode) - E°red (anode)
Trends
Spontaneity
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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
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ΔG = -nFE
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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
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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
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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
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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
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Portable, self-contained
electrochemical power
source
Batteries in series,
voltage is added.
Things to consider
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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
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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/fuelcell
8.swf
Corrosion
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Undesirable
spontaneous redox
reactions
Thin coating can
protect some metals
(like aluminum) –
forms a hydrated
oxide)
Iron - $$$$$
Protection
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Higher pH
Paint surface
Galvanize (zinc
coating) – why?
Zinc is a better anode
Called cathodic
protection – sacrificial
metal
More dramatic
Electrolysis
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Cells that use a battery or outside power
source to drive an electrochemical reaction in
reverse
Example NaCl → Na+ + ClReduction at the cathode, oxidation at the
anode
Voltage source pumps electrons to cathode.
Diagram
Solutions
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High temperatures necessary for previous
electrolysis (ionic solids have high MP)
Easier for solutions, but water must be considered
Example: NaF
Possible reductions are:
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Na+ + e- → Na(s) (Ered = -2.71 V)
2H2O + 2 e- → H2(g) + 2 OH- (Ered = -.83 V)
Far easier to reduce water!
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continue
Continued
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Look at possible oxidations:
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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
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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