Knowing Nernst: Non-equilibrium copper redox chemistry
Dec 21, 2015
Knowing Nernst:Non-equilibrium copper redox chemistry
Objectives:
(1) Calculate/measure stability of copper complexes
(2) Use ligands to change stabilities of metal species
HSAB concept: qualitative insights
Redox potentials/Nernst eqn: quantitative insights
Chemical species studies
• CuCl2• CuI
• Cu(NH3)42+
• Cu(en)22+
• Cu(salen)n+
• Charge vs oxidation state
Oxidation states
• Sum of oxidation states = ionic charge on species• Assumes unequal sharing of electrons
– more electronegative atom gets all of bond electrons
Oxidation states
• Sum of oxidation states = ionic charge on species• Assumes unequal sharing of electrons
– more electronegative atom gets all of bond electrons
• Examples: – MnO, MnO2, KMnO4
• What differences are found between compounds with difference oxidation numbers?
Atomic radius
Reactivity (redox potential)
Disproportionation
• 2 Fe4+ → Fe3+ + Fe5+
• 2 H2O2 → 2 H2O + O2
• 2 Cu+ → Cu0 + Cu2+
• Reverse of process: comproportionation
Sample redox potential calculation
CuCl2 + ammonia -> Cu(NH3)42+ + chloride
(1) Cu2+ + Iˉ + eˉ CuI 0.86V(2) Cu2+ + Clˉ + eˉ CuCl 0.54V
(3) I2 + 2eˉ 2Iˉ 0.54V (4) Cu+ (aq) + eˉ Cu(s) 0.52V(5) Cu2+(aq) + 2eˉ Cu(s) 0.37V(6) CuCl + eˉ Cu(s) + Clˉ 0.14V
(7) Cu(NH3)42+ + 2eˉ Cu(s) + 4NH3 -0.12V
(8) Cu2+(aq) + eˉ Cu+ (aq) -0.15V(9) CuI + eˉ Cu(s) + Iˉ -0.19V
(10) Cu(en)22+ + 2eˉ Cu + 2en -0.50V
Reduction: Cu2+(aq) + 2eˉ Cu(s) E0 = +0.37V (5)
Oxidation: Cu(s) + 4NH3 Cu(NH3)42+ + 2eˉ E0 = +0.12V (7*)
Net: Cu2+(aq) + 4NH3 Cu(NH3)42+ E0 = +0.49V
G0 = -nFE0
n = mol e-
F = 96,500 C / mol e-
E0 = standard reduction potential in V (1M conc, 1 atm pressure)
1 Joule = (1 Volt)(1 Coulomb)
Nernst Equation
n = number of mol e- R = 8.3145 J/K-mol
F = 96,500 C / mol e-
E0 = standard reduction potential in V (1M conc, 1 atm pressure)
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E0 =−RT
nFlnK eq
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E = E0 −RT
nFlnQ
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E = E 0 −0.0257V
nlnQ at 298 K
Hard vs. soft
• Describes the general bonding trends of chemical species (Lewis acids / Lewis bases)
• Hard acids prefer to bind to hard bases, while soft acids prefer to bind to soft bases
Hard: low polarizability, primarily ionic bonding
Soft: high polarizability, primarily covalent bonding
Lewis acids and bases
• Hard acids H+, Li+, Na+, K+ , Rb+, Cs+ Be2+, Mg2+, Ca2+ , Sr2+, Ba2+ BF3, Al 3+, Si 4+, BCl3 , AlCl3 Ti4+, Cr3+, Cr2+, Mn2+ Sc3+, La3+, Ce4+, Gd3+, Lu3+, Th4+, U4+, Ti4+, Zr4+, Hf4+, VO4+, Cr6+, Si4+, Sn4+
• Borderline acids Fe2+, Co2+, Ni2+ , Cu2+, Zn2+ Rh3+, Ir3+, Ru3+, Os2+ R3C+ , Sn2+, Pb2+ NO+, Sb3+, Bi3+ SO2
• Soft acids Tl+, Cu+, Ag+, Au+, Cd2+ Hg2+, Pd2+, Pt2+, M0, RHg+, Hg2
2+ BH3 CH2 HO+, RO+
• Hard bases F-, Cl- H2O, OH-, O2- CH3COO- , ROH, RO-, R2O NO3-, ClO4- CO3
2-, SO42- , PO4
3- NH3, RNH2 N2H4
• Borderline bases
Br- NO2-, N3-
SO32-
C6H5NH2, pyridine N2
• Soft bases H-, I- H2S, HS-, S2- , RSH, RS-, R2S
SCN- (bound through S), CN-, RNC, CO R3P, C2H4, C6H6 (RO)3P