Electrochemistry (Chapter 20)
Electrochemistry (Chapter 20) Electrochemical Cells -
Separate the half cell reactions , don’t allow them to physical come in contact with each other Beaker: one side – Cu anode, another side – Ag = cathode Connect wire through voltmeter / electrical thing from one side to the other Electrons need to travel from one side to the other Salt solution – ions of salt bridge filled with KNO3, K+ follows electrons, NO3= goes away from electrons, keeps electrical neutrality Two connected half cells The 2 half reactions are separated yet connected electrically (salt bridge, external wire) Solid metals = electrodes (anode, cathode) Movemen tof e= from anode to cathode generates voltage Movement of cation and anions thorough salt bridge maintains electrroneutrality in solution Oxidation (anode): surface of electrode erodes as Mn+ are produced, n e= then travel to cathode; M Mn+ + n eReduction (cathode): cathode gains mass as Mn+ ions gain n e= and form M(s); Mn+ + n e- M Oxidation: occurs at anode (Cu) Cu(s) --. Cu2+ (aq) + 2eReduction: occurs at cathode (Ag) 2Ag+(aq) + 2e- --. 2Ag(s) “ an ox and a red cat”, or vowels (O,A) / consonants (C, R) Overall cell reaction: Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s) Galvanic (voltanic) cells: result form spontaneous chemical reactions Electrolytic cell: use electricity to accomplish non-spontaneous chemical change sometimes an electrode is inert (does not chemically participate, but is needed for electron transfer to occur)
Standard Electrode Potentials -
absolute half-cell potentials cannot be measured all pontentials are measured relative to the standard hydrogen electrode (SHE) assigned a potential of 0V, E◦ = exactly 0 volt 2H+ (a=1) + 2e- h2 (g, 1 bar); simplify: a = 1 becomes [H+] = 1M; 1 bar = I atm Standard electrode potentials (E) report tendency for reduction to occur (Ered) Standard conditions: conc = 1M, pres
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iF E cell I pos – galvanic / spontaneous
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Ecell > 0 = spontaneous
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E cell = E red + E ox
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I clicker 3 ansewr – a
Reactions are spontaneous in one direction not spontaneous in the other
Spontaneity: Ecell and ΔG -
Electromotive force or cell potential (Ecell) Joule = volt x coulomb (or volt = energy/unit charge) Welec = zFEcell , where work = - deltaG ( pg 795) Z = #e- transferred; F = 96485 C/mol e- (faraday constant ΔG = -zFEcell ΔG◦ = -zFE◦cell Postitive Ecell or E◦cell = spontaneous Negative ΔG or ΔG◦= spontaneous ΔG = -RTlnK E◦cell = (RT/zF)inK K>1 means products are favourted K>1 means K>0 therefore E◦cell >0 Ecell – the higher the Ecell, the more likely a reaction goes towards completion ,more its shifted towards producsts in equilibrium Spontneous reaction has k>1 , E◦cell>0, ΔG◦ < 0
Non-standard Cell potentials (Ecell) -
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Concentrations do not equal 1M and/or pressures do not equal 1 atm Nernst equation : Ecell = E◦cell -(RT/zF)lnQ o R = gas constant (J/K mol), T = temp, z = # e- transferred, F = Faraday constant (96485 C/mol), Q = reaction quotient ( not at equilib, ratio of products to reactants at a given moment) Ecell = E◦cell =- (0.0592/z)logQ
Writing Q or K for thermodynamics -
When it’s a gas write it as a presur ein atm for equib eqn ICLICKER 4: decrease. Increase In [Fe3+] causes Q to decrease so lnQ decrease so Ecell increases All concentrations and pressures are taken with respect to reference concentrations (1M) or pressures (1bar = 1 atm) We can write a K expression with concentrations and pressures together because of activity (a) Under standard conditions, E◦cell = Ecell
Concentrations Cells -
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Identical electrodes with different ion concentraiotns Anode and cathode are the same: E◦cell = 0 Voltage still possible because of ion concentrations Connect two half cells made from the exact same solution and metal but change the concentration of either side of the half cell, low concentration in one high in the other – keeping things separate Uses electrons and transfers them from one half cell to the other half cell to make concentrations the same Cathode: 2H+ (aq, conc) + 2e H2(g) o We want concentration of H+ to go down so electrons goes to that side, reducing concentration of H+; reduction Anode: H2(g) 2H+ (aq, dil) + 2eo Want to take H2 gas to make more H+ because you want concentrations to be the same, give up 2 electrons; oxidation Overall: 2H+ (aq, conc) 2H+ (aq, dil ) No spontaneous reaction – negative Ecell
Separate the half cell reactions , don’t allow them to physical come in contact with each other Beaker: one side – Cu anode, another side – Ag = cathode Connect wire through voltmeter / electrical thing from one side to the other Electrons need to travel from one side to the other Salt solution – ions of salt bridge filled with KNO3, K+ follows electrons, NO3= goes away from electrons, keeps electrical neutrality Two connected half cells The 2 half reactions are separated yet connected electrically (salt bridge, external wire) Solid metals = electrodes (anode, cathode) Movemen tof e= from anode to cathode generates voltage Movement of cation and anions thorough salt bridge maintains electrroneutrality in solution Oxidation (anode): surface of electrode erodes as Mn+ are produced, n e= then travel to cathode; M Mn+ + n eReduction (cathode): cathode gains mass as Mn+ ions gain n e= and form M(s); Mn+ + n e- M Oxidation: occurs at anode (Cu) Cu(s) --. Cu2+ (aq) + 2eReduction: occurs at cathode (Ag) 2Ag+(aq) + 2e- --. 2Ag(s) “ an ox and a red cat”, or vowels (O,A) / consonants (C, R) Overall cell reaction: Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s) Galvanic (voltanic) cells: result form spontaneous chemical reactions Electrolytic cell: use electricity to accomplish non-spontaneous chemical change sometimes an electrode is inert (does not chemically participate, but is needed for electron transfer to occur)
Standard Electrode Potentials -
absolute half-cell potentials cannot be measured all pontentials are measured relative to the standard hydrogen electrode (SHE) assigned a potential of 0V, E◦ = exactly 0 volt 2H+ (a=1) + 2e- h2 (g, 1 bar); simplify: a = 1 becomes [H+] = 1M; 1 bar = I atm Standard electrode potentials (E) report tendency for reduction to occur (Ered) Standard conditions: conc = 1M, pres
-
iF E cell I pos – galvanic / spontaneous
-
Ecell > 0 = spontaneous
-
E cell = E red + E ox
-
I clicker 3 ansewr – a
Reactions are spontaneous in one direction not spontaneous in the other
Spontaneity: Ecell and ΔG -
Electromotive force or cell potential (Ecell) Joule = volt x coulomb (or volt = energy/unit charge) Welec = zFEcell , where work = - deltaG ( pg 795) Z = #e- transferred; F = 96485 C/mol e- (faraday constant ΔG = -zFEcell ΔG◦ = -zFE◦cell Postitive Ecell or E◦cell = spontaneous Negative ΔG or ΔG◦= spontaneous ΔG = -RTlnK E◦cell = (RT/zF)inK K>1 means products are favourted K>1 means K>0 therefore E◦cell >0 Ecell – the higher the Ecell, the more likely a reaction goes towards completion ,more its shifted towards producsts in equilibrium Spontneous reaction has k>1 , E◦cell>0, ΔG◦ < 0
Non-standard Cell potentials (Ecell) -
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Concentrations do not equal 1M and/or pressures do not equal 1 atm Nernst equation : Ecell = E◦cell -(RT/zF)lnQ o R = gas constant (J/K mol), T = temp, z = # e- transferred, F = Faraday constant (96485 C/mol), Q = reaction quotient ( not at equilib, ratio of products to reactants at a given moment) Ecell = E◦cell =- (0.0592/z)logQ
Writing Q or K for thermodynamics -
When it’s a gas write it as a presur ein atm for equib eqn ICLICKER 4: decrease. Increase In [Fe3+] causes Q to decrease so lnQ decrease so Ecell increases All concentrations and pressures are taken with respect to reference concentrations (1M) or pressures (1bar = 1 atm) We can write a K expression with concentrations and pressures together because of activity (a) Under standard conditions, E◦cell = Ecell
Concentrations Cells -
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-
-
Identical electrodes with different ion concentraiotns Anode and cathode are the same: E◦cell = 0 Voltage still possible because of ion concentrations Connect two half cells made from the exact same solution and metal but change the concentration of either side of the half cell, low concentration in one high in the other – keeping things separate Uses electrons and transfers them from one half cell to the other half cell to make concentrations the same Cathode: 2H+ (aq, conc) + 2e H2(g) o We want concentration of H+ to go down so electrons goes to that side, reducing concentration of H+; reduction Anode: H2(g) 2H+ (aq, dil) + 2eo Want to take H2 gas to make more H+ because you want concentrations to be the same, give up 2 electrons; oxidation Overall: 2H+ (aq, conc) 2H+ (aq, dil ) No spontaneous reaction – negative Ecell
