Zumdahl's Section 17.


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Electrochemistry: making charges work
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Slide 1

Zumdahl\'s Chapter 17 Electrochemistry: Making Charges Work

Slide 2

Galvanic Cells Cell Potential Std. Diminishment Potential, E ° Electrical Work Potential and Free Energy, G E \'s fixation reliance Nernst Equation K from E ° Batteries Corrosion Electrolysis Chapter Contents

Slide 3

Electrochemistry Conversion of concoction to electrical vitality (release). What\'s more, its opposite (electrolysis). Both subject to entropic alert: Convert reversibly to keep frameworks at harmony and change over all accessible substance work ( G) to and from the equal electrical work (QV). Electrons from REDOX responses.

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RedOx Half Reactions The e – are noticeable in ½ responses. 3 H 2 O 2  3 O 2 + 6 H + 6 e – 2 Au 3+ + 6 e –  2 Au 2 Au 3+ + 3 H 2 O 2  3 O 2 + 6 H + 2 Au But while ½ cells were a math comfort in stoichiometry, they are genuine in electrochemistry

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One ½ cell rxn. happens in every compartment. Zn  Zn 2+ + 2e – in the anode. Cu 2+ + 2e –  Cu in cathode. Be that as it may, not without an association. Cu Zn Galvanic Cells Cathode=Reduction Anode=Oxidation SO 4 2 – SO 4 2 – Zn 2+ Cu 2+ Zn + Cu 2+  Zn 2+ + Cu

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But even with an association of the terminals, no present streams. We have to permit lack of bias in the arrangements with a salt extension to move counterions. 2e – 2e – Ion ("salt") Bridge Cu Zn SO 4 2 – SO 4 2 – Zn 2+ Cu 2+ Zn + Cu 2+  Zn 2+ + Cu

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Standard Reduction Potentials, E ° The voltage created by the Zn/Cu galvanic cell is +1.1V under standard conditions . Standard conditions are: T = 25°C and P = 1 bar for gasses. Solids and fluids are unadulterated. Arrangements are 1 M in all species. E ° cell is entirety of ½ cell E ° values.

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½ cell Reduction Potentials All ½ cells are classified as decrease responses & doled out diminishment possibilities, E °. The lower diminishment potential ½ rxn is turned around to end up the oxidation. E ° oxidation = – E ° diminishment That makes unconstrained E ° cell > 0. Yet, E ° red can\'t be discovered w/o E ° bull !

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Origin for Reduction Potentials We had the same issue for S° particles and unraveled it by making H + uncommon. 2H + ( aq ) + 2e –  H 2 (1 bar) E °  0 V 1 bar H 2 streams over a Pt cathode, and the full E ° cell is doled out to the next anode. E ° SHE = 0 V. E.g., standard calomel cathode: Hg 2 Cl 2 ( s ) + 2e –  2 Hg( l ) + Cl – E ° SCE = +0.27V an all the more physically advantageous reference.

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Ag + e –  Ag .80V Cu 2+ + 2e –  Cu .34 2H + 2e –  H 2 .00 Fe 2+ + 2e –  Fe – .44 Zn 2+ + 2e –  Zn – .76 Mg 2+ + 2e –  Mg – 2.37 Etc. Keep in mind: turn around the lower potential to make it an oxidation rather than a lessening. A quick look at the standard lessening E °s at left explains to us why Cu is resistant to 1 M HCl while metals with lower E ° cheerfully rise off H 2 . Dynamic Metal Series

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Cu isn\'t undying H + doesn\'t do it. We fricasseed that penny not with HCl but rather with HNO 3 . So HNO 3 isn\'t just corrosive yet oxidizing corrosive! Cu 2+ + 2e –  Cu has E ° = +0.34V NO 3 – + 4H + 3e –  NO + H 2 O has E ° = + 0.96V So turning around the Cu and including HNO 3 gives a cell E ° = + 0.62 V Corroding Copper

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Galvanic Line Notation Shorthand for a complete redox cell is of the structure: Anode | anodic soln. || cathodic soln. | Cathode yet kept in touch with all on the same line. So making a cell of Cu consumption, Cu | Cu 2+ || NO 3 – , NO( g ), H + |Pt where all particles ought to be suffixed ( aq ) and both metals ought to have ( s ).

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Free Energy and Work Were all (aq) focuses in the Cu consumption cell at 1 M , the cell potential would be + 0.62V ( unconstrained ). Unconstrained responses have negative  G ° = max (non-PV) work. Electrical work = chargepotential n e moles of e – convey n e F Coulombs.   G ° = – n e F E ° J ( J = C V ) F = 96,485 C mol –1 , the Faraday const.

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Temperature Dependence of E ° Since E ° = –  G °/n F , and E = – G/n F so far as that is concerned, d E/d T = ( –1/n F ) d G/d T But d G = V d P – S d T , so d G/d T = – S Or d E/d T = +  S/n F   S °/n F where we\'ve assumed that neither S nor H will change much with moderate T. Since  S ° = + 124 J/mol K for an auto battery, it\'s harder to begin in winter. For 0°C, the 6 cell battery puts out 0.1V not exactly at 25°C

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Nernst Eqn: Potentials and Concentrations Both  G ° and E ° allude to unit (standard) fixations. In any case, at balance ,  G = 0 and the cell potential E = 0 too (see no ° ) .  G =  G ° + RT ln(Q)  – n e FE = – n e FE ° + RT ln(Q) E = E ° – 2.303 (RT/n e F ) log (Q) E = E ° – (59.1 mV/n e ) log(Q) @ 25°C

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K from E ° Just as  G =  G ° + R T ln(K) = 0 suggests  G ° = – R T ln(K), – n e F E ° = – R T ln(K) infers K = e + n e F E °/R T where, as some time recently, n e = moles of electrons required in the general response as composed! Large K can be computed.

Slide 17

Confession Time On slide 9, I touted the Hg 2 Cl 2/Hg couple as a helpful standard and drew its E ° from the table. Yet, S.C.E. remains for "immersed" calomel anode and E = 0.241 not E ° = 0.268 V (with soaked Cl – ) . Since Q = [Cl – ], by rearranging Nernst, we discover [Cl – ] sat\'d = 2.86 M . Cool.

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Potential from a SINGLE ½ Reaction?!? Don\'t we require an oxidation and also a lessening? Yes , yet they can be the same response (however for an inversion)! Focuses must contrast between the anode and cathode. I.e., Q should less be than 1 so log(Q) is negative; then despite the fact that E °=0 still E >0. The phone conveys Q to 1 at balance by evening out focuses in ½ cells.

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Ion-particular Electrodes [Ag + ] can be acquired by E from basic Ag wire alluded to SCE. [H + ] is a great deal more vital! pH cathodes, encased in glass, swap H + for Na + at silicate surface. Potential contrast along these lines actuated is adjusted for [H + ] outside . See your Harris § 15.4 for points of interest.

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Assaulted with Batteries "Battery" alludes to a progression of Galvanic cells whose E include. (Parallel hookup includes current, I, not E .) Rechargeable NiCad responses: Cd + 2 OH –  Cd(OH) 2 + 2e – NiO 2 + 2H 2 O + 2e –  Ni(OH) 2 + 2 OH – Notice the cancelation of OH – in conclusive response.  Q=1 generally so E altered! It doesn\'t rundown; it just stops.

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Better Batteries NiCad, however rechargeable, will acknowledge dynamically littler charges; "battery memory." NiMH replaces anode rxn with MH + OH –  M + H 2 O + e – with an any longer revive life. M may be Mg 2 Ni with  = 4.1 and successful H thickness twice H 2 (fluid)

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Best Battery OK, I\'m prejudiced. E ° 2 H 2 + 4 OH –  4 H 2 O + 4 e – +0.83V O 2 + 2 H 2 O + 4 e –  4 OH – +0.40V Is simply hydrogen burning; no Greenhouse gas. Best case of "power module" alleged on the grounds that H 2 and O 2 are not incorporated with the battery but rather supplied remotely. Notice that [OH – ] is again perpetual.

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Corrosion A battery is electrochemistry happening where you need it. Consumption is the place you don\'t . All M/MO x couples at E° < 0.4V are consumed even in burning arrangements: O 2 + 2 H 2 O + 4 e –  4 OH – E ° = 0.40 O 2 + 4 H + 4 e –  2 H 2 O E ° = 1.23 So corrosive improves. Q impact!

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Metal Corrosion Metal oxides are lower thickness (higher volume) than their metals. So oxide arrangement opens blooms of erosion and spreads. Salt shower is most exceedingly terrible; it\'s electrolytic! A few oxides (e.g., Cr 2 O 3 ) structure impenetrable oxide coats, moderating further O 2 assault.

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Sacrificial Anodes Mg  Mg 2+ + 2e – O 2 + 2H 2 O + 4e –  4 OH – Structural metals like Fe are flawlessly ensured by more dynamic (lower E °) metals like Mg. On the off chance that conductive contact is made, O 2 gets lessened (to H 2 O) on Fe by e – discharged from Mg. Supplanting the dynamic metal plate is less expensive than a rusted boat!

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Electromotive Force as a "Substance Reactant" If as opposed to doing work with a Galvanic cell potential, you supply an opposite potential, you run the response in the non - unconstrained bearing! Tough. Endo ergically. This is electrolysis , a blend. You supply E not e – ; the e – are taken from a cathode response, yet anode and cathode have swa pp ed .

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Electrolysis can just continue with a potential more negative than – E °. At that point the cell keeps running backward. Outside work supplies required  G . 2e – 2e – + Electrolysis Cell Cu Zn SO 4 2 – SO 4 2 – Zn 2+ Cu 2+ Zn + Cu 2+  Zn 2+ + Cu

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Electrolytic Stoichiometry Charge ( current  time = I  t ) decides measure of item. (Coulombs = Amperes  Seconds) Electrons are the constraining reactant in electrolysis. Moles electrons = n e = Q/F = I t/F The typical stoichiometric proportions change over between n e and moles of item.

Slide 29

Concentration Electrolysis? Does it bode well to run a focus cell in reverse? All you appear to do is to make a focus contrast as opposed to misusing one that tends to consistency. This is the way we purge metals ! Power debased metals to be anodes. They shed particles that are "plated" as immaculate metal on the cathodes!

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Making Active Metals You can\'t "plate" Na , say, out of a watery arrangement! It will essentially redox respond with H 2 O to make NaOH(aq). We electroly

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