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Ch 10 Lecture 1 Bonding Basics. Evidence of Electronic Structure What is Electronic Structure? Electronic Structure = what orbitals electrons reside in and their energies Coordination Compounds The residence of electrons in the metal s, p, and d orbitals

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ch 10 lecture 1 bonding basics
Ch 10 Lecture 1 Bonding Basics
  • Evidence of Electronic Structure
    • What is Electronic Structure?
      • Electronic Structure = what orbitals electrons reside in and their energies
      • Coordination Compounds
        • The residence of electrons in the metal s, p, and d orbitals
        • The complete MO description of the entire complex
      • Bonding theories for transition metal complexes must explain characteristics of the compound influenced by d-electrons
      • Certain experimental data must match the theorized electronic structure
    • Thermodynamic Data
      • Stability (or Formation) Constants
        • Indicators of Complex Stability
slide2
Equilibrium constants of complex formation

Cu2+ + NH3 [Cu(NH3)2+]

[Cu(NH3)]2+ + NH3 [Cu(NH3)2]2+

[Cu(NH3)2]2+ + NH3 [Cu(NH3)3]2+

[Cu(NH3)3]2+ + NH3 [Cu(NH3)4]2+

  • Combining the individual steps gives the overall stability constant b4

b4 = K1 x K2 x K3 x K4 = 6.8 x 1012

  • Large b values indicate favorable reactions and/or stable complexes
  • Small b values indicate unfavorable reactions and/or unstable complexes
slide3
Can be measured by pH titrations, UV-Vis titrations, etc…
  • Thermodynamic Data
    • DG = -RTlnK = DH – TDS allows calculation of free energy, entropy, and enthalpy of a reaction from stability constants
    • These values can help explain bonding in coordination compounds, but detailed theoretical calculations are needed to fully rationalize a given set of data
    • Most useful to see differences in a single parameter in a series of related compounds
    • Example: DH might decrease from Br- to Cl- to F- for their complexation to Fe3+ (Hard-Hard Interaction)
    • This data is not very useful in predicting structure or bonding interactions of coordination compounds in general
slide4
Magnetic Susceptibility
    • Magnetic Properties reveal numbers of unpaired electrons
    • Hund’s Rule requires maximum number of unpaired electrons in degenerate orbitals
    • Magnetic Descriptions
      • Diamagnetic = all paired electrons = slightly repelled by magnetic field
      • Paramagnetic = unpaired electrons = strongly attracted by magnetic field
    • Magnetic properties are determined by one of several experimental methods

not

slide5
Molar Magnetic Susceptibility = cM = cm3/mol = data from an experimental determination
  • Magnetic Moment = Calculated value from cM and theoretical treatment of magnetism. m = 2.828(cMT)½Bohr magnetons
  • Theoretical basis of magnetism
    • Electron Spin: a spinning charged particle would generate a spin magnetic moment = mS (not really spinning but a property of an electron)
    • mS = + ½ or – ½ only because the electron can either “spin” clockwise or counter clockwise (up or down arrows)
    • S = Spin Quantum Number = S mS
      • Example: S = 3(+ ½ ) = 3/2
      • Example: S = +1/2
slide6
Angular momentum quantum number = l = describes the orbital shape

s = 0, p = 1, d = 2, f = 3

  • mL = Orbital Angular Momentum = + l, l -1,…..- l
    • Each number is assigned to one orbital of the set
    • Example: p-orbitals
  • L = Orbital Quantum Number = S mL
    • Example: L = +1 + 0 + -1 = 0
    • Example: L = 2(+1) + 0 = 2
  • Magnetic moment m depends on both S and L

mS+L = g[S(S+1) + ¼ L(L+1)]½

g = gyromagnetic ratio ~ 2.00 for an electron

  • The Orbital contribution is small for first row transition metals, so we can use a spin only magnetic moment = mS = g[S(S+1)]½ = [n(n+2)]½

n = number of unpaired electrons

+1 0 -1

slide7
Example: calculate mS for high spin Fe3+
    • Determine d-electron count by counting back three positions on the periodic table (3+) and counting how far from left on the periodic table you are: d5
    • Arrange the d-electrons in the 5 d-orbitals as high spin
    • Apply mS = g[S(S+1)]½ = 2[5/2(7/2)]½ = 5.92 Bohr magnetons

Or [n(n+2)] ½ = [5(7)] ½ = (35) ½ = 5.92

slide8
Valence Bond Theory
    • History
      • Proposed by Pauling in the 1930’s
      • Describes bonding using hybrid orbitals filled with e- pairs
      • Extension of Lewis/VSEPR to include d-orbitals
    • Theory
      • Metal ions utilize d-orbitals in hybrids
      • Octahedral complexes require 6 hybrid orbitals
        • d2sp3 hybridization of metal Atomic Orbitals provides new MO
        • Ligand lone pairs fill the hybrid orbitals to produce the bond
        • d-orbitals can come from 3d (low spin) or 4d (high spin)

Fe3+

Co2+

slide9
Problems with the theory
        • High energy 4d orbitals are unlikely participants in bonding
        • Doesn’t explain electronic spectra of transition metal complexes
  • Crystal Field Theory
    • History
      • Developed to describe metal ions in solid state crystals only
      • M+ is surrounded by A- “point charges”
      • Energies of the d-orbitals are “split” due to unequal geometric interactions with the point charges
      • Does not take into account covalency and molecular orbitals
      • Has been extended to do so in Ligand Field Theory
    • Theory
      • Place degenerate set of 5 d-orbitals into an octahedral field of (-) charges (L:)
      • The electrons in the d-orbitals are repelled by the (-) charge of the ligands
      • The dz2 and dx2-y2 orbitals are most effected because their lobes point directly along x,y,z axes where the point charges are
      • The dxy, dxz, and dyz orbitals aren’t destabilized as much
slide10
The energy difference between these orbital sets is called “delta octahedral” = Do
    • The low energy set has t2g symmetry and are stabilized by –0.4 Do each
    • The high energy set has eg symmetry and are destabilized by +0.6 Do each
    • The total energy of the 5 d-orbitals is the same as in the uniform field = 0

(2)(+0.6 Do) + (3)(-0.4 Do) = 0

slide11
CFSE = Crystal Field Stabilization Energy = how much energy is gained by the electrons in the 5 d-orbitals due to their splitting
    • Co(III) = d6 low spin

(6e-)(-0.4 Do) = -2.4 Do stabilization

    • Cu(II) = d9

(6e-)(-0.4 Do) + (3e-)(+0.6 Do) = -0.6 Do stabilization

    • Cu(I) = d10

(6e-)(-0.4 Do) + (4e-)(+0.6 Do) = 0 Do stabilization