Unit 3 Electrochemistry/Battery Book Notes
18.1) Half-Reactions
Half-reaction: either oxidation or reduction process (but doesn’t happen alone, both
happen together)
18.2) The Half-Reaction Method of Balancing Redox Equations
-
When putting oxidation + reduction reactions together, make sure same # of e on both
sides (multiply by coefficient if needed)
o Everything that’s lost must be gained b/c nothing’s created/destroyed
o 1) Separate redox equation into 2 half-reaction equations
o 2) Balance # of atoms in half-equations except O and H (do these last)
o 3) Balance electric charge by adding # of e needed to make that side neutral/get
same net charge (gained on left; lost on right)
o 4) Adjust coefficients so each half-reaction has same # of e -
o 5) Add 2 half-reactions (gather all reactants on reactants sides, products on
products side.
o 6) Simplify overall reaction if necessary (add/divide coefficients)
Redox Reactions (Rxns) in Acidic Solution (sol’n)
o H and H O in2olved
+
o Balance O w/ H O an2 H w/ H
Redox Rxns in Basic sol’n
- +
o OH must appear instead of H
o Balance equation as if it’s in acidic sol’n and then add as many OH as there are
+
H they cancel and make water, and leave some hydroxide 18.3) A Qualitative Description of Voltaic Cells
Voltaic cell: device that uses spontaneous redox reaction to produce electricity
Electrode: metal strip used in electrochemical experiment
o Electrode equilibrium: established when electrode’s immersed in sol’n of its ions
Higher up on activity series more easily oxidized (that direction in equation’s strongly
favored)
Half-cell: consists of an electrode immersed in sol’n of ions
Salt bridge: inverted U-shaped tube containing salt sol’n; porus plugs @ ends of tube
prevent all sol’n from flowing, but allow ions to migrate; prevents buildup of charges
(anions move through it to neutralize any cation buildup, & move out too)
Metals connected by wire to voltmeter
Current can’t move in sol’n form b/c oxidized/reduced @ electrode
o cations & anions carry electric charge
More likely oxidized cation goes into salt bridge anion from salt bridge comes into
cation’s sol’n/// anion leaves less likely oxidized cation’s sol’n cation of salt bridge
comes into sol’n (p. 756)
Some Important Electrochemical Terms
o Electrochemical cell: device that combines 2 half cells w/ appropriate connections
btwn electrodes + sol’ns
o Anode: electrode where oxidation occurs; source of e in voltaic cell negative
electrode (but really, e buildup = so small here so not much difference in net
-
charge @ 2 electrodes); potential energy of e (electric potential) greater here than
@ cathode o Cathode: electrode where reduction occurs; positive electrode, receives e ; -
o *Mnemonic device (CAR, AUTO CathodeAReduction, AnodeUTOxication)
o Flow Anode cathode
o Electric potential measured in volts; Electric charge measured in coulomb (C)
1V = 1J/C
o Voltmeter measures difference (diff) in electric potential btwn 2 pts in a circuit
If the 2 pts = electrodes, the potential diff (cell potentialcell) = cell
voltage) = driving force that propels e from anode cathode
o Overall redox rxn = cell rxn
Cell Diagrams
o Anode = on left
o Cathode = on right
o Single vertical line (|) represents boundary btwn diff phases, such as btwn
electrode & a sol’n
o Double vertical line (||) represents salt bridge/other porous barrier separating 2
half cells
o Sometimes can use electrode that doesn’t participate in redox (inert); just
furnishes surface on which an electrical potential = established.
18.4) Standard Electrode Potentials
Can know voltage w/o using specific measurements find potential difference
Standard hydrogen electrode (SHE): hydrogen gas @ 1 bar pressure bubbled over inert
platinum electrode & into aqueous sol’n in which concentration’s been adjusted so that
activity of H3O (H ) is exactly a = 1. + +
o a H = 1 [H ] = 1M
o P H 2= 1 bar P H2=1 atm
Standard electrode potential, E°, for any half-reaction = based on tendency for reduction
to occur @ an electrode; measured w/ all sol’n species present @ unit activity (a = 1),
which = about 1M, and gases @ bar pressure, which = about 1 atm
o When no other metal’s indicated as electrode material, potential is that for
equilibrium established on an inert surface, such as platinum metal
o SEP for SHE = arbitrarily set @ exactly 0 volts
To find SEPs, set up diagram
o Standard cell potential (E° cellE°(cathode) – E°(anode) b/c travels from anode
to cathode; get values from table
o E° cell E°(right) – E°(left)
o Positive voltage means right ions = more readily reduced than left ones, vice versa
for negative
o Anode’s gotta be connected to negative terminal of voltmeter, cathode to +, if
reversed, reading = negative
Electrode potentials & cell potentials = intensive properties magnitudes are fixed,
concentrations specified do not depend on size of half-cell/voltaic cell
Cell potentials can be calculated for redox reactions even if voltaic cells aren’t involved
18.5) Electrode Potentials, Spontaneous Change, and Equilibrium
Total work done = (cell voltage)(#moles of electrons transferred btwn electrodes,
n)(electric charge per mole of electrons (Faraday constant,F, = 96485 coulombs per
mole)) o ω elec nFE ;cell units volt * coulomb
o Max amt of useful work system can do = – ΔG; max amt of work that can be
realized as electrical work
o – ΔG = ω = nFE ;
elec cell
o *w/o degree sign E celleans electrodes’ conditions aren’t standard conditions;
can put degree sign on G too, means same thing, for now, it’s nonstandard
conditions; for standard, put degree symbol on both sides
Criteria for Spontaneous Change in Redox Rxns
o If E = positive, then ΔG = negative & rxn in forward direction = spontaneous
cell
o If E cell negative, then ΔG = positive & rxn in forward direction is
nonspontaneous
o If E cell 0, system’s @ equilibrium
o When cell rxn’s reversed, E cell ΔG change signs
Even if cell terminology used, any predictions apply either rxns = carried out in voltaic
cells or by mixing reactants
The Activity Series of the Metals Revisite
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