3 Electrochemistry Battery_Book Notes.docx

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CHEM 1061
Hyunjoo Im

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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