Download Donor Interactions Donor Interactions 10.4.1 Sigma 10.4.1

10.4 Angular Overlap
- an approach to bonding that is useful for making estimates of E of orbitals in coordination complexes
- estimate the strength of interaction b/w ligand orbitals & metal d orbitals based on the overlap
- Why this term ?? :
amount of overlap depends strongly on the angular arrangement of the M orbitals & the angles of ligand
10.4.1 Sigma
Sigma--Donor Interactions
- Strongest interaction: b/w M-dz2 & ligand-p
reference interaction for other σ-interaction
bonding orbital: larger ligand component
decreased MO E by eσ
antibonding orbital: larger M component
increased MO E by eσ
Fig.10.21
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10.4.1 Sigma –Donor Interactions
- Table 10.10: changes in orbital E due to other interactions b/w M d orbitals & ligand orbitals
need to justify the # qualitatively
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10.4.1 Sigma –Donor Interactions
- Example on p.390) [M(NH3)6]n+ , octahedral
only σ-interaction (NH3 → no π orbital)
donor orbital on NH3 → pz
→ px, py: used in bonding w/ H
- calculation of the orbital E in a complex,,
1) d orbital: ∑ of # for the appropriate ligands in the
vertical column
2) ligand orbital: ∑ of # for all d orbitals in the
horizontal row
Fig.5.31
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10.4.1 Sigma –Donor Interactions
- Metal d orbitals - 1) dz2: strongest w/ 1 & 6 – raise E by eσ
: weak w/ 2, 3, 4, 5 – raise E by 1/4 eσ
increase by total 3 eσ
2) dx2-y2: position 1, 6 – no interaction
position 2, 3, 4, 5 – raise E by 3/4 eσ
increased by total 3 eσ
3) dxy, dxz, dyz: no interaction w/ ligands
remains unchanged
- Ligand orbitals - 1) ligand 1, 6 w/ dz2: lowered by eσ
w/ other d: no interaction
2) ligands 2, 3, 4, 5 w/ dz2: lowered by 1/4 eσ
w/ dx2-y2: lowered by 3/4 eσ
∴ Each ligand orbital is lowered by eσ !!
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10.4.1 Sigma –Donor Interactions
- Fig.10.22: resulting E pattern
decribe how the M complex is stabilized
(X 2) d orbitals of M: increase E
(X 3) d orbitals of M: remains unchanged
(X 6) ligand orbitals: lower E
* net stabilization: 12 eσ for the bonding pair
Fig.10.22
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10.4.2 Pi
Pi--Acceptor Interactions
- CO, CN-, PR3 (phosphine): π acceptors w/ empty orbitals
- Strongest π interaction: b/w dxz of M & π* of ligand
- Fig.10.23:
- eπ < eσ → ∵ π overlap is weaker than σ-overlap
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10.4.2 Pi
Pi--Acceptor Interactions
- Table 10.11: Pi interaction
1) M d orbital dz2, dx2-y2: no π interaction at all
(1-6)
dxy, dxy, dyz: stabilized by 4eπ
2) ligand orbitals → raise E
- d e- → occupy the bonding MO w/ - 4eπ
- Fig. 10.24: σ-donor & π-acceptor ligands
∆o = 3eσ + 4eπ
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10.4.2 Pi
Pi--Acceptor Interactions
- Example on p. 394) [M(CN)6]n-
M’s dxy, dxz, dyz lowered by 4eπ
X 6 ligands → increased by 2eπ
∆o (t2g - eg split) = 3eσ + 4eπ
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10.4.3 Pi
Pi--Donor Interactions
- Interaction b/w occupied ligand p, d or π-orbitals
&
metal d-orbitals
- Similar to π-acceptor
π acceptor case except the reversed E b/w M & ligands
M d orbitals: E ↑
ligand π-orbitals: E ↓
- Fig.10.26:
- Fig.10.25:
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10.4.3 Pi
Pi--Donor Interactions
- Example on p.395) [MX6]nhalide ion: σ-interaction → donate e- via py
: π-interaction → donate e- via px, pz
both σ- & π-donor
dz2 & dx2-y2 orbitals: no π-interaction
no effect on the E of these d orbitals
dxy, dxz and dyz orbitals: π-interaction w/ four ligands
dxy → w/ 2, 3, 4, 5 → by total 4eπ ↑
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10.4.4 Types of Ligands and the Spectrochemical Series
- Ligands can be classified by their donor & acceptor capabilities
- NH3: σ-donors only, no π-interactions
simple bonding (Fig.10.4)
- ∆ (ligand field split): depends on the relative E of M & ligand
the degree of overlap
- ∆ (en) > ∆ (NH3)
strongest effect
= order of proton basicity
- Ligand field strength of halide ions
F- > Cl- > Br- > I= order of proton basicity
- Ligands
Li
d w/
/ occupied
i d p orbitals:
bit l π-donors
d
( σ-bonding
(+
b di e-)
∆ ↓ (section 10.4.3)
most halide complexes: high-spin configuration
- ex) the order of other examples:
H2O > F- > RCO2- > OH- > Cl- > Br- > IInorganic Chemistry 2
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10.4.4 Types of Ligands and the Spectrochemical Series
- Ligands w/ π* or d orbitals: π back-bonding is possible
becomes π acceptors
∆↑
ex)) CO,
CO CN- > phenanthroline
h
th li > NO2- > NCS- When these lists are combined,,
spectrochemical series (from strong π-acceptor to strong π-donor)
CO, CN- > phen > NO2- > en > NH3 > NCS- > H2O > F- > RCO2- > OH- > Cl- > Br- > ILow spin
High spin
Strong field
Weak field
Large ∆
donor only
Small ∆
π acceptors
π donors
Strong interactions w/ TM’s orbitals
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10.4.5 Magnitudes of eσ, eπ, and ∆
▪ Charge on metal
- Changing the ligands or the M → affects the magnitudes of eσ, eπ, ∆
result in a change in the # unpaired e- ex) H2O: weak field ligand
w/ Co2+ in Oh → [Co(H2O)6]2+, high-spin, 3 unpaired ew/ Co3+ in Oh → [Co(H2O)6]3+, low-spin, no unpaired e-
Fig.10.27
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10.4.5 Magnitudes of eσ, eπ, and ∆
▪ Different Ligands
- ex) [Fe(H2O)6]3+ vs [Fe(CN)6]3+
high-spin
low-spin
balance among ∆, Πc, Πe → determine high- or low-spin
- Tetrahedral complex: ∆t is small
low-spin tetrahedral complex is unlikely!!
instead,, low-spin octahedral complex!!
- Table 10.12, Table 10.13: angular overlap parameters
observed trends 1) eσ > eπ
∵ σ interaction: more direct orbital overlap b/w nuclei
π interaction: small overlap, no direct toward each other
2) σ , π ↓ as size ↑
∵ bond length ↑ → overlap ↓
as EN of halide ion ↓
∵ the pull ↓
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10.4.5 Magnitudes of eσ, eπ, and ∆
- Table 10.12: ligands – listed in the spectrochemical series order
ex) for Cr3+ → CN- is listed first
highest in the spectrochemical series
∵
π-acceptor (eπ is ‘-’)
en
→e
→ NH3
: affected only by eπ values (σ-donor ability)
→ halide ions: ∵ at the bottom of the series
π-donor & σ-donor
▪ Special Cases
- The angular overlap model: can describe the electronic E of complexes
w/ diff. shapes or
w/ combinations of diff.
diff ligands
: estimate the magnitudes of eg & eπ w/ diff. ligands
- ex) [Co(NH3)4Cl2]+: low-spin
∴ magnetic property does not depend on ∆o
but, ∆o effect on the visible spectrum (Chapter 11)
- Helpful to compare E of diff. geometries
ex) for 4-coordinate complex → tetrahedral or sqaure-planar?? (section 10.6)
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10.5 The John
John--Teller Effect
- There can not be unequal occupation of orbitals w/ identical E.
molecules distorted
no longer degenerate orbitals!!
ex) Cu(Ⅱ), d9, Oh - Fig.10.28
3 e- in the two eg w/o J-T effect
J-T effect → slight shape change of a complex
change in E of the orbitals
distortion 1) elongation (along one axis)
or
2) compression (along one axis)
→ in Oh,, eg*: direct toward ligands
larger effect on E level
t2g: No direct overlap
small effect
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10.5 The John
John--Teller Effect
- strong J-T effect: when eg* are unequally occupied
weak J-T effect: when t2g are unequally occupied
- Expected J-T effect
- ex) significant J-T effects, Cr(Ⅱ) (d4)
high-spin Mn(Ⅲ) (d4)
Cu(Ⅱ) (d9)
Ni(Ⅲ)
( ) ((d7)
low-spin Co(Ⅱ) (d7)
- low-spin Cr(Ⅱ): distorted from Oh → to D4h
: 2 absorption bands
one in Vis
one in near IR
: if pure Oh → only one d-d transition (Chapter 11)
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10.5 The John
John--Teller Effect
- ex) - Cu(Ⅱ): significant J-T effects → elongation
: affect equilibrium constant
: [Cu(NH3)4]2+ → distorted Oh w/ 2 H2O at greater distance
→ putting the 5th & 6th NH3 is difficult
see the formation constant
the bond distance of two axial position: longer!!
smaller equilibrium constant for these positions!!
∴ Cu(Ⅱ) complexes prefer square-planar geometry!!
[Cu(H2O)6]2+ + NH3
[Cu(NH3)(H2O)5
]2+
[Cu(NH3)(H2O)5]2+ + H2O
+ NH3
[Cu(NH3)2(H2O)4]2+
[Cu(NH3)3(H2O)3]2+
[Cu(NH3)4(H2O)2]2+
+ NH3
K1 = 20,000
+ H2O
K2 = 4,000
[Cu(NH3)3(H2O)3
]2+
+ H2O
K3 = 1,000
+ NH3
[Cu(NH3)4(H2O)2
]2+
+ H2O
K4 =
200
+ NH3
[Cu(NH3)5(H2O)]2+
K5 =
0.3
[Cu(NH3)(H2O)]2+ + NH3
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[Cu(NH3)2(H2O)4]2+
[Cu(NH3)6]2+ + H2O
2011 Fall
+ H2O
K6 = very small
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