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CH7. Intro to Coordination Compounds

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Inner-sphere vs outer-sphere

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Nomenclature 1. Learn common ligand names (Table 7.1)

Ex: :OH2 aqua :O2− oxo (oxido) :CN− cyano (cyanido) :Br− bromo (bromido) :NH3 ammine

Note that anionic ligands end in “o”

2. List ligands in alphabetical order

3. Metal name at end, add “ate” if it’s an anionic complex

some common names – ferrate, stannate, plumbate, cuprate

4. Add (and metal oxidation number in Roman numerals)

or add metal (and total complex charge in Arabic numerals)

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Nomenclature ex: [Cu(OH2)6]2+ is hexaaquacopper(II) or hexaaquacopper(2+)

[CuCl4]− is tetrachlorocuprate(III) or tetrachloridocuprate(III)

5. Add prefixes to indicate number of each ligand type

mono, di, tri, tetra, penta, hexa

or use bis, tris, tetrakis if less confusing due to ligand name

ex: [PtBr2{P(CH3)3}2 ] is dibromobis(trimethylphosphine)platinum(II)

Stereoisomers cis- and trans-platin. The cis isomer is an anti-cancer drug. ~ C2v ~D2h

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Cis-platin binding to DNA

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Nomenclature 6. To write the formula: [metal, then anionic ligands, then neutral ligands] net charge superscript

7. Special ligands:

a. ambidentate -SCN (thicyanato) vs −NCS (isothiocyanato)

− NO2 (nitrito) vs −ONO (isonitrito)

[Pt(SCN) 4 ] 2 − D 4h tetrathiocyanatoplatinate(II )

[Cr(NCS)(NH3)5] 2+ pentaammineisothiocyanatochromium(III )

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Nomenclature b. bidentate – ligands bind to M at two sites

ex: H2NCH2CH2NH2 ethylenediamine (en)

[Cr(en)3]3+ tris(ethylenediamine)chromium(III)

View looking down C3 axis

D3 (-> no σ, no S axes, chiral)

enantiomers

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Nomenclature Another bidentate example is acetato

c. polydentate ligands – bind at multiple sites

ex: tetraazamacrocycles

porphine (a simple porphyrin)

the 4 N atoms are approximately square planar

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Geometric Isomers There have distinct physical and chemical properties

Oh coordination MX5Y 1 isomer

MX4Y2 2 isomers (cis or trans)

MX3Y3 2 isomers (fac = C3V or mer = C2V )

ex: [CoCl2(NH3)4]+ tetraamminedichlorocobalt(III)

cis – purple trans – green

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Optical Isomers

Enantiomers = non-superimposable mirror images of a chiral molecule

enantiomers have identical physical properties (except in a chiral environment, for example retention times on a chiral column are not the same)

enantiomers rotate the plane of polarized light in opposite directions (optical isomers)

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Polymetallic complexes (also called cage compounds)

no direct M-M bonding

ex:

S8 + NaSR + FeCl3 → [Fe4S4(SR)4]n− model for ferrodoxins

MeOH (dry) / N2

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Cluster compounds direct M-M bonding

ex: [Re2Cl8]2− octachlorodirhenate(III)

D4h (eclipsed)

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Crystal Field Theory Oh complexes – put 6 e− pairs around central metal in Oh geometry

this splits the 4 d-orbitals into 2 symmetry sets

t2g (xz, yz, xy) and eg (x2 – y2, z2)

∆0 can be determined from spectroscopic data (see Table 8.3)

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UV/Vis spectrum for Ti(OH2)63+

20,300 cm-1 (wavenumber units)

= 493 nm (wavelength units) = 243 kJ/mol (energy units)

violet solution

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Crystal Field Theory ∆0 depends on:

1. ligand (spectrochemical series)

∆0 I− < Br− < Cl− < F− < OH− < NH3 < CN− < CO

weak field strong field

more complete list in text

2. metal ion

∆0 greater for higher oxidation number – stronger, shorter M-L interaction

∆0 greater going down a group – more diffuse d-orbitals interact more strongly with ligands

∆0 Mn2+ < Fe2+ < Fe3+ < Ru3+ < Pd4+ < Pt4+ small ∆ large ∆

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Ligand Field Stabilization Energy for electronic config t2g

x egy the LFSE = (0.4x − 0.6y) ∆0

high spin case

# d electrons 0 1 2 3 4 5 6 7 8 9 10

e− config - t2g1 t2g

2 t2g3 t2g

3eg1 t2g

3eg2 t4eg

2 t5eg2 t2g

6eg2 t2g

6eg3 t2g

6eg4

LFSE (∆0) 0 0.4 0.8 1.2 0.6 0 0.4 0.8 1.2 0.6 0

# unpaired e− 0 1 2 3 4 5 4 3 2 1 0

depends of relative values of ∆0 and pairing energy.

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High spin vs low spin d4

t2g3eg

1 t2g4

LSFE = 0.6 ∆0 LFSE = 1.6 ∆0 − PE

high spin low spin

(weak field) (strong field)

[Cr(OH2)6]2+ [Cr(CN)6]4−

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∆Ηhyd for first-row TM2+ ions

All are high spin complexes

M2+(g) → H2O → [M(OH2)6]2+ (aq)

∆H calc from Born Haber analyses

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Magnetic Measurements Magnetic moment (µ) is the attractive force towards a magnetic field (H)

µ ≈ √N(N + 2) µB ↑

where N = number of unpaired electrons

N µ/µB 1 1.73 2 2.83 3 3.87 4 4.90 5 5.92

this is the paramagnetic contribution from unpaired e− spin only, it ignores both spin-orbit coupling and diamagnetic contributions

ex: [Mn(NCS)6]4− experimental µ/µB = 6.06,

Mn(II) is d5 it must be a high spin complex

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CN = 5

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d-orbital splitting in a Td field

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CFT for CN 4 For Td complexes

∆T << ∆0 due to fewer ligands and the geometry of field vs ligands

ex: Δ

[CoCl4] 2− 3300 cm −1 [Co(OH2)6]3+ 20,700 cm −1

therefore Td complexes are nearly always high spin (pairing E more important than LFSE)

Co(II) d7 LSFE = 1.2∆T

ex: Fe3O4 magnetite Fe(II)Fe(III)2O4

oxide is a weak field ligand, so high spin case

Fe(II) is d6 (only in Oh sites); Fe(III) is d5 (1/2 in Oh sites, ½ in Td sites)

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Square planar complexes

D4h is a common structure for d8 complexes (full z2, empty x2 – y2

orbitals)

Group 9: Rh(I), Ir(I)

Group 10: Pt(II), Pd(II)

Group 11: Au(III), for example AuCl4−

Note: [Ni(CN)4]2− is D4h but [NiCl4]2− is Td

Ni(II) has a smaller ∆ than Pd, PT so Td is common

but we see D4h with strong field ligands

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Tetragonal distortion of Oh

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Jahn-Teller effect Jahn-Teller effect: degenerate electronic ground states generate structural disorder to decrease E

Ex: [Cu(OH2)6]2+ Cu(II) d9

We see a tetragonal distortion

But fluxional above 20K, so appears Oh by NMR in aqueous solution

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Jahn-Teller effect

CuF2

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Ligand Field Theory

CFT does not explain ligand field strengths; MO theory can

Start with SALCs that are ligand combinations shown to the right

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MO for Oh TM complexes

SF6 - no metal d valence orbitals considered

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π-bonding in Oh complexes

π-acceptor ligands Increase ∆O Example: CO

π-donor ligands Decrease ∆O Example: Cl-

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Oh character table