Knowing Nernst:Non-equilibrium copper redox chemistry

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

Kstability = [AB] / [A][B]

softerharder

most stable complexes

least stable complexes

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 

Experimental Details

--Part G: watch out for oil drips and ethanol flames

--do not throw away stir bars--recover them

--dissolve all of the H2salen and Cusalen--no precipitates