8
CHAPTER
BONDING IN TRANSITION
METAL COMPOUNDS AND
COORDINATION COMPLEXES
8.1 Chemistry of the Transition Metals
8.2 Introduction to Coordination Chemistry
8.3 Structures of Coordination Complexes
8.4 Crystal Field Theory: Optical and Magnetic
Properties
8.5 Optical Properties and the Spectrochemical
Series
8.6 Bonding in Coordination Complexes
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Emerald
3BeO∙Al2O3∙
6SiO2
with some
Al3+ replaced
by Cr3+
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8.1 CHEMISTRY OF THE TRANSITION
METALS
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- Decreasing radii for small Z transition atoms
→ Increase in Zeff
- Increasing radii for large Z transition atoms
→ Increase in electron-electron repulsion
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 Lanthanide contraction: bad shielding by 4f orbitals
→ the radii of the 6th period ~ the 5th period
→ decrease in atomic and ionic radii by increasing Z
along the 6th period
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 melting point: function of
the bond strength in solids
- roughly correlated with the
number of unpaired e-
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 Enthalpy of hydration of M2+ ions
M2+(g) → M2+(aq): Hhyd
= Hof(M2+(aq)) – Hof(M2+(g))
 Lowering of Hhyd from a line
→ due to crystal field stabilization
 Anomalies of Mn
→ due to the stable half-filled d shell
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 Oxidation states
more common oxidation state
 Increasing tendency toward higher oxidation states among
heavier transition elements in the same group:
Fe (2,3) → Ru (2,3,4,6,8), Ni(2,3) → Pd(2,4)
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 Hard and Soft Acids and Bases
 Pearson (1963)
~ Extension of Lewis’ definition –
electron pair acceptor (acid) and donor
(base) – by adding categories ‘hard’
and ‘soft.’
~ 'Hard' species: small, high charge
states, low electronegativities, weakly
polarizable
~ 'Soft' species: large, low charge
states, high electronegativities, strongly
Ralph Pearson
(US, 1919 - )
polarizable
~ ‘Borderline’ species
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 Prediction of chemical reactivities of inorganic reactions
~ Preferred direction: hard acid/hard base or soft acid/soft
base
HgF2(g) + BeI2(g) → BeF2(g) + HgI2(g)
s/h
h/s
h/h
s/s
AgBr(s) + I–(aq) → AgI(s) + Br–(aq)
s/b
s
s/s
b
EXAMPLE 8.2
Predict whether the following reactions will occur.
(a) CaF2(s) + CdI2(s) → CaI2(s) + CdF2(s)
(b)Cr(CN)2(s) + Cd(OH)2(s) → Cd(CN)2(s) + Cr(OH)2(s)
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NO
YES
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8.2 INTRODUCTION TO COORDINATION
CHEMISTRY
 Formation of Coordination Complexes
 Werner’s investigation:
Compound 1: CoCl36NH3 (orange-yellow)
Compound 2: CoCl35NH3 (purple)
Compound 3: CoCl34NH3 (green)
Compound 4: CoCl33NH3 (green)
Alfred Werner
(Swiss,1866-1919)
Nobel prize in
chemistry(’13)
 Treatment with HCl → did not remove NH3
AgNO3 + Cl- → AgCl(s) in the ratio of 3 : 2 : 1 : 0
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 Conductivity measurements:
Compound 1: [Co(NH3)6]3+(Cl–)3
→ Conductivity of Al(NO3)3
Compound 2: [Co(NH3)5Cl]2+(Cl–)2 → Conductivity of Mg(NO3)2
Compound 3: [Co(NH3)4Cl2]+(Cl–) → Conductivity of NaNO3
Compound 4: [Co(NH3)3Cl3]
→ Nonelectrolyte
→ Concept of “coordination sphere”
around the central metal ion
inner and outer coordination sphere
→ Formation of an octahedral complex
In the above complexes, NH3 and Cl- that are
attached to Co are called LIGANDS
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CuSO4∙5H2O
anhydrous CuSO4
→ [Cu(OH2)4]SO4∙H2O
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 Monodentate ligands
mono “one” and dens “tooth”
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 Bidentate ligands
(‘ox’)
(‘en’)
 Chelating ligands: chelate (G. chele, “claw”)
[Pt(en)3]4+ ~ 3 bidentates
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[Co(EDTA)]– ~ 1 hexadentate
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 Naming coordination compounds
1) Single word for a coordination complex ~ [prefix-ligand-metal]
2) Cation first followed by anion ~ K[…] or […]Cl
3) Ending with the suffix “-o” for anionic ligand, chlorido (Cl),
no change for neutral ligands except aqua (H2O), ammine (NH3),
carbonyl (CO).
Note: “chloro” for Cl in a compound ligand
4) Prefixes for the number of ligands ~ di-, tri-, tetra-, penta-, hexa-, …
(bis-, tris-, tetrakis-, … for ligands with di- (etc) in their names)
5) Alphabetical ordering for many ligands
6) Roman numeral (oxidation state) in (..) after the name of metal
~ […cobalt(III)]Cl
or
K[…ferrate(III)]
anionic complex ions: the ending “-ate”
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 Ligand substitution reactions
[Ni(OH2)6]2+(aq) + 6 NH3(aq) → [Ni(NH3)6]2+(aq) + 6 H2O
Another example
_
HCl(aq)
NH 3(aq)
2+
Cu(H2O)6 (aq)
CuCl4(aq)
Green
Pale blue
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Cu(NH3)62+(aq)
Deep blue
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Difference between ‘inert’ and ‘labile’
 ‘Inert’ coordination complex:
stable (inert)
[Co(NH3 )6 ]3 (aq)  6H3O (aq)  [Co(H2O)6 ]3  6NH4 (aq)
Energy
thermodynamically unstable, kinetically
takes a week
Reaction
thermodynamically unstable, kinetically
unstable (labile)
[Co(NH3 )6 ]2 (aq)  6H3O (aq)  [Co(H2O)6 ]2  6NH4 (aq)
Energy
 ‘Labile’ coordination complex:
takes a matter of seconds
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Reaction
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8.3 STRUCTURES OF COORDINATION COMPLEXES
Octahedral complexes with geometrical isomers
(complexes of type MA2B4 (or MA2B2; B is bidentate)
cis-[Co(NH3)4Cl2]+
trans-[Co(NH3)4Cl2]+
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cis-[CoCl2(en)2]+
trans-[CoCl2(en)2]+
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 Octahedral complexes with mer / fac isomers
(Complexes of type MA3B3)
mer-isomer: Similar ligands define a meridian
of the octahedron
fac-isomer: Similar ligands define a face of an octohedron
-- all three groups are 90° apart.
mer-Co(NH3)3(Cl)3
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fac-Co(NH3)3(Cl)3
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 Tetrahedral complexes
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~ Dominant for four-coordinate complexes
~ No geometrical isomers for tetrahedral
complexes of MA2B2
 Square planar complexes
~ Au3+, Ir+, Rh+, Ni2+, Pd2+, Pt2+
~ cis-[Pt(NH3)2Cl2]
(anticancer drug, ‘cisplatin’)
~ trans-[Pt(NH3)2Cl2]
 Linear geometry
~ Ions with d10 configuration: Cu+, Ag+, Au+,
Hg2+
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 Chiral Structures
 Optical isomers are molecules that rotate plane polarized light
 Enantiomers (Gk. eάτιος, “opposite”, and μέρος, “part or
portion”) are optical isomers whose structures are non-
superimposable mirror images (they lack reflection-rotation
symmetry)
 Chiral center (chirality [G. χειρ (kheir), "hand"] ~ handedness) is
a central atom around which enantiomers are formed
 A racemic mixture has equal amount of enantiomers (net
rotation of plane polarized light = 0)
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Octahedral
complexes of
type MA3
(A is bidentate)
E.g. enantiomers of
the [Pt(en)3]4+ ion
Octahedral
complexes of
type MA2B2C2
E.g. enantiomers of
all-cis [Co(NH3)2(H2O)2Cl2]+
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 EDTA (ethylenediaminetetraacetate) ion
Hexadentate ligand, sequestering metal ions
Antidote for lead poisoning, preserves freshness of oil
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8.4 CRYSTAL FIELD THEORY: OPTICAL AND
MAGNETIC PROPERTIES
 Crystal Field Theory
~ Ionic description of metal-ligand bonds
~ Ligands are treated as point charges approaching
the central metal ion
Octahedral coordination complexes
 Degeneracy of d-orbitals lifted into two groups :
d
z2
, d x2  y 2
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
and
d
xy
, d yz , d z 
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 Crystal Field Theory
• Ligands such as a halide or oxide are regarded as an electrostatic,
point charge, or point dipole type, which set up an electrostatic field.
A
B
Cr3+
metal d orbitals
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spherical
charges
octahedral
environment
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o = crystal
field
splitting
energy
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Fig. 8.17 An octahedral crystal
field increases the energies of
all five d orbitals, but the increase
is greater for the dz2 and dx2 - y2
orbitals.
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• Electron configuration of octahedral complexes d1-d3
by Hund’s rule
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-From d4 to d7 octahedral complexes there are two possibilities,
illustrated for d4 (E.g. Mn(III) complexes)
If o is large (strong-field ligands), t2g4 has a lower energy.
: low-spin complex, minimum number of unpaired eIf o is small (weak-field ligands), t2g3eg1 has a lower energy.
: high-spin complex, maximum number of unpaired eeg
eg
E
t2g
t2g
Low spin (t2g4) configuration
e-e repulsion
low-spin configuration
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High spin (t2g3eg1) configuration
ligand-ligand repulsion
high-spin configuration
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- Example:
d4
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octahedral complexes of Mn(III)
d z2
Mn(CN)63LOW SPIN
Mn(H2O)63+
HIGH SPIN
3
o
5
eg
3
5 o
o
d x 2- y 2
dz 2
eg
o
2 o
5
2 o
5
5 x degenerate
d orbitals (3d4)
d x2- y 2
dxy
d yz
d xy
dxz
t 2g
Weak field configuration
H2O weak field ligand
d yz
d xz
t 2g
5 x degenerate
d orbitals (3d4)
Strong field configuration
CN– strong field ligand
Fig. 8.18. Electron configuration for (a) high spin (large o) and (b) low spin
(small o) octahedral crystal field splitting energies for Mn(III) complexes
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 Crystal Field Stabilization Energy (CFSE)
The amount by which the (otherwise equal) energy levels for the
d electrons of a metal ion are split by the electrostatic field of the
surrounding ligands in a coordination complex.
 Energy difference between electrons in an octahedral crystal
field and those in the hypothetical spherical crystal field.
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 Square planar crystal field
sp > 1.6 0
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 Tetrahedral crystal field
t = 4/9 o
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Fig. 8.20. Correlation diagram showing the relationships among
d-orbital energy levels in crystal fields of different symmetries.
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 Magnetic properties
 Magnetic susceptibility
~ Strength of a sample’s interaction with a magnetic field
 Paramagnetic compounds
~ One or more unpaired electrons
~ Large, positive magnetic susceptibility
~ Attracted by the magnetic field
→ “weigh” more
~ Prevalent among transition-metal complexes
 Diamagnetic compounds
~ All of the electrons are paired
~ Small, negative susceptibility
~ Repelled by the magnetic field
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8.5 OPTICAL PROPERTIES AND THE
SPECTROCHEMICAL SERIES
 Transition-metal complexes
~ absorb visible light → colorful
E.g. [Co(NH3)5Cl]2+ ion absorbs greenish yellow light (~530 nm)
Only red and blue light transmitted
→ purple (complementary color)
 Wavelength of the strongest absorption, max
E  h , so o  h  hc / max
d10 complex ~ colorless (no absorption, all d-levels are filled)
High-spin d5 complex ~ weak absorption (spin flip required)
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Cr(CO)6
[Co(NH3)5(OH2)]Cl3
K3[Fe(C2O4)3]
K3[Fe(CN)6]
[Co(en)3]I3
Colors of the hexaaqua complexes of metal ions prepared from their nitrate salts.
E.g. [Co(H2O)6]2+
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 Spectrochemical series
~ An ordering of ligands according to their ability to cause
crystal field splittings.
 Spectrochemical series for ligands
I  Br   Cl  F ,OH  H2O  : NCS  NH3  en  CO,CN
Weak-field ligands (high spin)
Intermediate-field ligands
Strong-field ligands (low spin)
 Spectrochemical series for metal ions
Mn2+ < Ni2+ < Co2+ < Fe2+ < Fe3+ < Co3+ < Mn4+ < Pd4+ < Ir3+ < Pt4+
 Crystal field theory cannot explain the spectrochemical series!
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8.6 BONDING IN COORDINATION
COMPLEXES
 Valence bond theory
 dsp3 hybrid orbitals
~ linear combination of one s, three p atomic orbitals
and the dz2 atomic orbital
~ five equivalent new hybrid orbitals
~ trigonal bipyramid, PF5, CuCl5–
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 d2sp3 hybrid orbitals
~ linear combination of one s,
three p atomic orbitals
and dz2, dx2-y2 orbitals
~ six new hybrid orbitals
~ octahedron, SF6
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 Molecular orbital theory
 Ligand field theory
~ Failure of CFT and VB theories to explain the spectrochemical
series
~ MO description for ligands
 Construction of s MOs for octahedral complexes (of 1st row
D-block metals)
~ Interaction between the metal 4s orbital with six ligand orbitals
→ ss and ss* orbitals
~ Interaction between three metal p orbitals with three ligand orbitals
→ triply degenerate sp and sp* orbitals
~ Interaction of the dz2 and dx2-y2 orbitals with ligand orbitals
→ a pair of sd and sd* orbitals
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Fig. 8.27. Formation of s bonding MOs
from overlap of metal and ligand orbitals.
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Antibonding MOs
MO correlation diagram
for octahedral Cr(III)
complex ([CrCl6]3-): s
bonding only
Nonbonding MOs
Bonding MOs
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 Formation of  and * bonds
(1) Interaction between an empty metal d orbital with a filled atomic
ligand p orbital. E.g. 3p orbitals of Cl–
(2) Interaction between a filled metal d orbital with an empty ligand *
antibonding molecular orbital. E.g. CO, CN–
→ metal-to-ligand (M-L)  donation or  backbonding
-  and * MOs:
M d orbital - L p orbital
or M d orbital - L * orbital
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(3) Overlap of each of the metal nonbonding dxy, dyz,
and dxz orbitals with four ligand p orbitals
→ Formation of three pairs of
bonding and antibonding
MOs, t2g and t2g*.
Fig. 8.30. Bonding  MO by constructive overlap
of a metal dxy orbital with four ligand p orbitals.
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 Order of bonding strengths for different ligands
 Weak-field ligands (small o)
→ Overlap between occupied p() bonding orbitals of
ligands (Br–, Cl–, CO) with t2g orbitals of metal
→ Increase in energy of t2g and decrease in o
 Strong-field ligands (large o)
→ Overlap between unoccupied * antibonding orbitals of
ligands (CO, CN–) with t2g orbitals of metal
→ Lowering of energy of t2g orbitals by  back-bonding (M→L)
 Intermediate-field ligands ~ H2O, NH3
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t2g*
Empty ligand
p (*) orbitals
eg*
eg
eg
Empty ligand
p (*) orbitals
eg*
t2g*
E
Partially filled
metal d orbitals
Filled ligand
p () orbitals
Partially filled
metal d orbitals
acceptor (M
ligands
t2g
(a)
donor (M
ligands
t2g
L)
L)
Filled ligand
p () orbitals
(b)
Fig. 8.31. (a) (ML) [or (b) (ML)]  donation showing a reduction
(or increase) in Δo compared with that from s bonding alone.
(a) Slight increase in energy of t2g electrons (in t2g* MOs)
(b) Significant lowering in energy of t2g electrons
due to  back-bonding → Electrons of t2g MOs are delocalised
into unoccupied *(L)
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 Summary of the MO picture (Ligand Field Theory)
of bonding in octahedral coordination complexes
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IIlustrated
for V2+,Cr3+,Mn4+
(d3)
Cl-, Br- ligands
e.g. [CrCl6]3–
CO, CN–, NO+
Ligands e.g.
Mn(CN)4
H2O, NH3 ligands
e.g. [V(H2O)6]2+
Fig. 8.32. Effect of  bonding on the energy-level
structure for octahedral coordination complexes.
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10 Problem Sets
For Chapter 8,
2, 8, 18, 26, 32, 44, 46, 58, 64, 66
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