Transition Metal Coordination Chemistry
What is a transition metal?
Prof S.M.Draper
2.05 SNIAMS Building
smdraper@tcd.ie
Recommended books
M.J. Winter, d-block Chemistry, Oxford Chemistry Primers, OUP, 2001
M.S. Silberberg, Chemistry, 3rd Ed, McGrawHill, 2003 (chapter 23)
C.E. Housecroft, A.G. Sharpe, Inorganic Chemistry, 1st Ed, PrenticeHall, 2001
J.E. Huheey, E.A. Keiter, R.L. Keiter, Inorganic Chemistry, 4th Ed., HarperCollins, 1993
1
Electrons in atoms
z
Shapes of d-orbitals
4s
y
x
3d
4p
Sc
Ti
z
z
z
V
y
Cr
y
y
Mn
x
x
x
Fe
Co
yz
xz
xy
x2-y2
z2
Ni
Cu
Zn
2
Working out numbers of d-electrons from oxidation states:
How many d-electrons has the metal?
1 2
3 4 5 6
7 8
9 10 11 12
1st: how many electrons are there in the shell?
O
- count along the periodic table
e.g. Mn = 7 electrons
ox =
Cu = 11 electrons
en =
O
-O
H2N
O-
NH2
2nd: how many electrons are lost?
- oxidation state
complex
O.S. of L
O.S. of M
no. d electrons
e.g. Mn(VII) = 7 electrons lost Cu(II) = 2 electrons lost
[Cr2O7]2-
-2
+6
d0
[MnO4]-
-2
+7
d0
0
+3
d1
-1, 0
+2
d8
-2
+3
d5
[Ag(NH3)2]+
3rd: how many electrons left over?
- subtract
e.g. Mn(VII) =
[Ti(H2O)6]3+
Cu(II) =
[Co(en)3]3+
[PtCl2(NH3)2]
[V(CN)6]4-
Hence the only valance electrons available in a transition metal ion are d-electrons
[Fe(ox)3]3-
3
Colour of transition metal complexes
biological
activity
magnetic behaviour
colour
Ruby
Corundum
Al2O3 with
impurities
geometry
What’s interesting about
Sapphire
Transition Metal Complexes??
Al2O3 with
coordination
number
medical
applications
oxidation states
octahedral metal centre
Corundum
and
impurities
coordination number 6
Emerald
Beryl
AlSiO3 containing Be with
impurities
4
Haemoglobin
Oxygen carrier in blood
O2
Porphyrin-Fe transition metal complex
Cisplatin
[PtCl2(NH3)2]
Fe(II) ion is octahedrally coordinated
NN
NN
Fe
NN Fe
NN
2C
HOHO
2C
Coordination number 6
NR
NR
HO C
2
HO 2 C
square planar Pt(II)
coordination number 4
cis-isomer
the first of a series of platinum coordination complex-based anti-cancer drugs
(Platinol-AQ)
5
Alfred Werner - Nobel Prizewinner 1913
[Co(NH3)6]Cl3
[Co(NH3)5Cl]Cl2
3+
CoCl3 . 6NH3
yellow
xs Ag+
3 moles AgCl
CoCl3 . 5NH3
purple
xs Ag+
2 moles AgCl
CoCl3 . 4NH3
green
xs Ag+
1 mole AgCl
xs Ag+
0 moles AgCl
CoCl3 . 3NH3
[Co(NH3)4Cl2]Cl
+
2+
Werner's conclusions
1.
The metal is in a particular
(primary valancy)
2.
The complex has a fixed
(secondary valancy)
3.
The ligands are bound to the metal via a bond which resembles a covalent bond
6
What is a coordination complex?
Lewis Acid-Base Concept: parallel with main group
chemistry
• Adducts formed by the reaction of metal ions (Lewis acids) with
several ligands (Lewis bases).
H
Ni2+
+
H 2O
H
+
6 :NH 3
X+/-
2+
O
. .:
H 3 N:
NH 3
n
Ni
NH 3
Ni 2+
n+/-
Ni
NH 3
2+
:NH 3
Central metal ion or atom surrounded by a set of ligands
NH 3
The ligand donates two electrons to the d-orbitals around the metal forming a
• Properties of products differ from those of separate reactants (i.e.,
colour, composition, magnetic properties, etc.).
1
3
7
F
F
zz
F
B
z
z
F
F
B
z
z
F
z
z
H
F
B
F
H
N
H
zz
H
F
N
H
H
NH3
_
H N > BF
BF3
3
zz
6
L
H
N
H
L
L
3+
zz
3
zz
NH3
e.g. [Co(NH3)6]3+
z
z
+
zz
z
z
zz
H
H
z
z
z
z
zz
N
z
z
H
z
z
zz
z
z
H
+
Co3+
H3N
H3N
z
z
z
z
z
z
NH3
z
z
NH3
zz
L
L
NH3
L
= coordinate or dative bond
1
C
Group 14
O
CO
carbon monoxide
e.g. NH3, H2O, OH-, CO32-
CO, PPh3, C2H4, SRH, CN-, SCN-
Small donor atoms
Electronegative
H
H
Group 15
N
NH3
ammonia
O
H2O
water
H
PPh3
phosphine
S
SR2
thioether
Not very polarisable
Less electronegative
Easily polarisable
R
H
Group 16
P
Larger donor atoms
H
R
e.g. Fe(III), Mn(II), Cr(III)
e.g. Ag(I), Cu(I)
Small metals (1st row)
Larger metals (2nd + 3rd row)
High oxidation state
Low oxidation state
2
π - bonded ligands
Group 14
C
N
z
z
CNcyanide
-
Phphenyl
CH2
Cl
K+
Cl
H2C
CH2
RC
CR
Pt
Cl CH2
Group 15
N
O
z
z
Group 16
O
Group 17
H
X
NOnitrous
N
OHhydroxide
S
halide
C
C
H
S
NCS-
z
z
N
isocyanate
z
z
[PtCl3(η2-C2H4)]-
SCNthiocyanate
hydride
Monodentate
donor atom per ligand
Bidentate
donor atoms per ligand
Tridentate
three
donor atoms per ligand
Multidentate
many
donor atoms per ligand
3
Anionic bidentate ligands
Neutral bidentate ligands: 2 donor atoms
O
H 2N
NH2
zz
zz
e.g. [PtCl2(en)]
π-donor bidentate ligand
O
O
O
oxalate = ox2-
-
Fe
C
C
O
O
O
O
1,2-diaminoethane = ethylene diamine = en
[Fe(CO)3(η4-C4H6)]
acetate = ac-
H 3C
C
O
Ph2P
zz
PPh 2
zz
N
zz
N
zz
N
zz
zz
2,2'-bipyridine
bpy
1,10-phenanthroline
phen
R'
O
N
O
O
Pd
R'
1,2-diphenylphosphineethane
dppe
OH
R
N
M
O
N
HO
R
C N
H
N C
H
Pd(II)-oxime complex
4
Tetradentate ligands
Tridentate ligands: three donor atoms
Hexadentate ligand
tetraanion of
H 2N
zz
zz
NH2
N
NH2
NH
zz
N
zz
zz
ethylenediaminetetraacetic acid
N
N
NH2
zz
EDTA
NH2
O
tris(2-aminoethyl)amine
diethylenetriamine
2,2':6',2"-terpyridine
dien
tpy
O
-
tren
O
N
1,2,4-triazacyclonane
HN
NH
zz
NH
N
HN
N
O
N
O
O
-
O
N
N
N
zz
macrocyclic ligand
NH
N
O
N
HN
N
zz
N
H
M
porphyrin
phthalocyanin
5
Inner coordination sphere =
Outer coordination sphere =
H2O
SO42-
H2O
H2O
Mn2+
OH2
OH2
H2O
[Mn(OH2)6][SO4]: outer sphere complex
H2O
OSO3
H2O
Mn2+
OH2
H2O
H2O
[Mn(OH2)5(SO4)5]: inner sphere complex
6
Ligand Exchange Reactions
[Fe(OH2)6]2+
2+
[Fe(OH2)5(CN)]+
+
H2O
+
+
hexaaquo iron(II)
complex
CN-
+
cyanide
ion
+
K1
=
const
[H2O]
=
[Fe(OH2)5(CN)+] [H2O]
[Fe(OH2)62+] [CN-]
1
The reaction continues….
The reaction continues….
[Fe(OH2)5(CN)]+
+
K2
=
CN-
[Fe(OH2)4(CN)2]+
+
H2O
[Fe(OH2)4(CN)2] + CN-
[Fe(OH2)3(CN)3]-
+
H2O
β3 =
[Fe(OH2)3(CN)3]- + CN-
[Fe(OH2)2(CN)4]2- +
H2O
β4 =
[Fe(OH2)2(CN)4]2- + CN-
[Fe(OH2)(CN)5]3-
+
H2O
β5 =
[Fe(CN)6]4-
+
H2O
β6 =
[Fe(OH2)4(CN)2]
[Fe(OH2)5(CN)+] [CN-]
K1K2
=
β2
[Fe(OH2)5(CN)+]
=
[Fe(OH2)62+] [CN-]
x
[Fe(OH2)3(CN)3-]
[Fe(OH2)62+] [CN-]3
[Fe(OH2)2(CN)42-]
[Fe(OH2)62+] [CN-]4
[Fe(OH2)(CN)53-]
[Fe(OH2)62+] [CN-]5
[Fe(OH2)4(CN)2]
[Fe(OH2)5(CN)+] [CN-]
[Fe(OH2)(CN)5]
+ CN-
Overall Stability Constant
βn
=
[Fe(CN)64-]
[Fe(OH2)62+] [CN-]6
[MLn]
[M] [L]n
Stability constants are often expressed in log form ie log βn
2
Values of βn
The Chelate Effect
2+
4-
+ 6
+ 6
[M(NH3)6]2+
n+
NH3
H3N
M
NH3
NH3
H3N
NH3
β6
=
[Fe(CN)64-]
[Fe(OH2)62+] [CN-]6
log β6 = 35
n+
H2N
H2
N
N
H2
[M(en)3]2+
NH2
M
NH2
H2N
~ 1035
[M(OH2)6]n+
+
6 NH3
[M(NH3)6]n+ +
6 H2O
[M(OH2)6]n+
+
3 en
[M(en)3]n+
6 H2O
+
Kn < … < K3 < K2 < K1
3
ΔGo = ΔHo - TΔSo
Entropy of chelate formation
ΔGo = -RT ln K
[Cu(OH2)6]2+ +
β2107.7
ΔGo = ΔHo - TΔSo
Enthalphy changes similar
ΔHo
= - 46 kJ
+
[M(OH2)6]n+
+
mol-1
6 NH3
[M(NH3)6]n+
+
6 H2O
3 en
[M(en)3]n+
+
6 H2O
2 H2O
+
2 H2O
log β2 = 7.7
ΔSo = - 8.4 J K-1mol-1
[Cu(OH2)4(en)]2+
en
β2 1010.6 log β1 = 10.6
ΔHo = - 54 kJ mol-1
ΔSo is large and positive
+
Entropy changes differ
[Cu(OH2)6]2+ +
[M(OH2)6]n+
[Cu(OH2)4(NH3)2]2+
2 NH3
ΔSo = + 23 J K-1mol-1
- TΔSo is large and negative
4
Thermodynamic vs Kinetic Stability
The Macrocyclic Effect
Complexes containing macrocyclic rings have enhanced stability compared to acyclic
ligands
2+
H
N
N
H
2+
H
N
H
Cu
Cu
N
H2
N
N
H2
log K1 = 23.9
acyclic chelating ligand
H
N
N
H
log K1 = 28.0
macrocyclic ligand
e.g
[Cr(OH2)6]3+
d3
kinetically inert
slow substitution of Ls
cf.
[Fe(OH2)6]3+
hs d5
kinetically labile
rapid substitution of Ls
(thermodynamic stability similar)
ΔG always favours formation of macrocyclic complex
5
Tour of Coordination Numbers and their Geometries
Most Common Geometries of Transition Metal Complexes
The Kepert Model
The shape of a complex is usually dictated by the number of coordinated atoms.
The coordinated atoms are as far away from each other as possible.
Tetrahedral
109o 28'
C.N. 4
Square Planar
90o
C.N. 4
Trigonal bipyramidal
120o + 90o
C.N. 5
Square based pyramidal
90o
C.N. 5
Octahedral
90o
C.N. 6
Non-
bonding electrons are ignored.
Note:
•
regular geometries are often distorted
•
structural features of multinuclear complexes are described in terms used
for individual metal centres
•
when energy differences between different structures are small, fluxional
behaviour may be observed
6
Coordination number 2
Coordination number 5
Cu(I), Ag(I), Au(I), Hg(II)
[CuCl2]-
trigonal bipyramid
[Au(CN)2]-
square-based pyramid
180o
90o
Coordination number 3
90o
120o
[Cu(CN)2]-
[HgI3]CN
CN
Cu
Cu
C
N
N
C
C
120o
N
N
C
Cu
Cu
CN
CN
The two structures have very similar energy
n
7
trigonal bipyramid
3-
Cl
Cl
Cu
square-based pyramid
Cl
NC
Cl
CN
CN
[Co(CN)5]3-
[CuCl5]3-
+
N
Co
Co
NC
Cl
N
3-
CN
O
N
N
Br
[Co(Me6tren)Br]+
O
O
V
O
O
[VO(acac)2]
8
Tetrahedral complexes
Favoured by steric requirements
optical isomerism
1
1
2
4
3
[CoCl4
4
3
]2-
[MnO4]90o
2
non-superimposable mirror images
[NiCl4]2-
109o
1
Square Planar Geometry
Square planar geometry
e.g. [PtCl4]2[AuBr4][Co(CN)4]2-
cisplatin
Square planar complexes are formed by
i.e. group 10
Ni2+, Pd2+, Pt2+
Au3+
cis-[PtCl2(NH3)2]
trans-[PtCl2(NH3)2]
cis-diamminedichloroplatinum(II)
trans-diamminedichloroplatinum(II)
Square planar complexes are formed by
e.g. CN-
2
Coordination number 6
Octahedral geometry
Geometrical Isomerism
[ML4X2]
e.g. [Mn(OH2)6]2+
[Cr(CO)6]
Staggared ligands vs. Eclipsed ligands
C+
trans-[Co(NH3)4Cl2]+
cis-[Co(NH3)4Cl2]+
green
violet
C+
do metals
e.g. WMe6
3
Octahedral geometry
[ML3X3]
Geometrical Isomerism
Octahedral geometry
[M(Κ2L)2X2]
Geometrical and Optical Isomerism
e.g. [Co(en)2Cl2]+
trans geometrical isomer
H
N
Cl
H
N
Co
N
H
Cl
N
H
mer-[Co(NH3)3(NO2)3]
cis geometrical isomer
fac-[Co(NH3)3(NO2)3]
two optical isomers
Non-superimposable mirror images = Δ and Λ enantiomers
4
Octahedral geometry
Optical Isomerism
Tetragonal Distortion:
Tetragonal elongation
z-axis
2+
[M(Κ2L)3]
2 long axial & 4 shorter
equatorial bonds
N
N
Ru
N
N
N
N
[Cu(OH2)6]2+, [Cr(OH2)6]2+
Tetragonal compression z-axis
M
[Ru(bpy)3]2+
2 short axial & 4 longer
equatorial bonds
trans-[Co(NH3)4(CN)2]+
* Small distortion but symmetry lower than octahedral.
5
Isomerism
Trigonal Distortion:
• Two faces on opposite sides of octahedron move either towards
or away from each other
N
N
Rotate
by 60°
Elongation
N
Stereoisomerism
Isomers have the same empirical
Isomers have the same M-L bonds,
formula but the atoms connectivities
but the atoms are arranged differently
differ
in space
Ionisation isomerism
Geometrical isomerism
M
N
N
Complexes of tacn can show this type of distortion [M(tacn)2]2+
N
Hydration isomerism
cis/trans
Coordination isomerism
mer/fac
Linkage isomerism
N
N
Structural Isomerism
N
tacn = 1,3,7-triazacyclononane
Polymerisation isomerism
Optical isomerism
Δ and Λ
6
Ionisation isomerism
Coordination isomerism
Exchange of a ligated anion with a counterion
no ppt
[Co(NH3)5Br]SO4
Ag+
[Co(NH3)5(SO4)]Br
AgBr
Ba2+
Ba2+
[Cu(NH3)4][PtCl4]
[Pt(NH3)4][CuCl4]
square planar
[Co(NH3)6][Cr(CN)6]
[Cr(NH3)6][Co(CN)6]
octahedral
BaSO4
no ppt
Polymerisation isomerism
Solvate isomerism
Exchange of a neutral ligand for an anionic ligand
e.g.
[Co(OH2)6]Cl3
e.g.
[Pt(NH3)4][PtCl4]
[Pt(NH3)2Cl2]
violet
[Co(OH2)5Cl]Cl2
light green
[Co(OH2)4Cl2]Cl
Both polymers have the empirical formula [Pt(NH3)2Cl2]n
dark green
7
Higher Coordination Numbers
Linkage isomerism
Coordination Number 7
hν
[Co(NH3)5(NO2)]2+
Δ
[Co(NH3)5(NO2)]2+
yellow
red
nitro-complex
nitrito-complex
[Pd(NCS)2(PPh3)2]
isocyanate
[Pd(SCN)2(PPh3)2]
thiocyanate
[WBr3(CO)4)](distorted)
[TaF7]2[ZrF7]3-
8
Coordination Number 8
Na3[Mo(CN)8]
(nBu4N)3[Mo(CN)8]
Coordination Number 9
[ReH9]2-
9
Symmetrical ligand field
z
Shapes of d-orbitals
orthogonal
z
y
x
M
x and y
z
z
z
y
y
y
x
x
x
x2-y2 yz
z2
xz xy
Energy
yz
xz
xy
x2-y2
z2
x2-y2 yz
z2
xz xy
1
Effect of ligand field on metal d-orbitals
The difference in energy between the eg and the t2g energy levels is the
. This is given the value 10 Dq.
symmetrical field
octahedral ligand field
eg
Δo or 10 Dq
t2g
2
Electron configurations: d1 ion
3+
The nature of Δoct
[Ti(OH2)6]3+
e.g. [Ti(OH2)6]3+
490-580 nm
violet solution in water
one d-electron in a t2g orbital
the complex has a Crystal Field Stabilisation Energy (CFSE) of - 0.4 Δoct
Absorption spectrum: λmax = 510 nm = 243 kJ mol-1
3
d4 ions
e.g. [V(OH2)6]3+
e.g. [Cr(OH2)6]3+
+ 0.6 Δoct
CFSE =
- 0.4 Δoct
= - 0.6 Δoct
+ 0.6 Δoct
CFSE =
- 0.4 Δoct
= - 1.6 Δoct + P
4
Factors affecting the magnitude of Δoct
High and Low spin Complexes
High spin complex
Low spin complex
e.g.
Δ is small
Δ is large
electrons occupy eg and t2g orbitals
electrons pair in t2g oribtals before
singly before pairing
occupying eg orbitals
[Fe(OH2)6]2+
Δoct = 10 000 cm-1
[Fe(OH2)6]3+
Δoct = 14 000 cm-1
[Co(OH2)6]2+
Δoct = 9 700 cm-1
[Co(OH2)6]3+
Δoct = 18 000 cm-1
e.g.
[Co(NH3)6]3+
Δoct = 22 900 cm-1
[Rh(NH3)6]3+
Δoct = 34 100 cm-1
[Ir(NH3)6]3+
Δoct = 41 000 cm-1
5
eg
eg
I- < Br- < S2- < SCN- ≈ Cl-< NO3- < F- < OH- < ox2t2g
< H2O < NCS- < CH3CN < NH3 ≈ en < bpy
< phen ≈ NO2- < PR3 < CN- ≈ CO
t2g
M.J. Winter, d-block Chemistry, p.83
6
Octahedral Crystal Field
ions, Oh field
d-orbital energies in…
free ion
spherical ligand field
+ 0.6 ∆oct
octahedral ligand field
- 0.4 ∆oct
x2-y2 yz
z2 xz
xy
+ 0.6 ∆oct
x2-y2 yz
z2 xz
- 0.4 ∆oct
xy
the complex has a Crystal Field Stabilisation Energy (CFSE)
ions, Oh field
ions, Oh field
High Spin
eg
eg
+ 0.6 ∆oct
+ 0.6 ∆oct
t2g
t2g
- 0.4 ∆oct
- 0.4 ∆oct
Low Spin
eg
+ 0.6 ∆oct
t2g
- 0.4 ∆oct
+ 0.6 ∆oct
- 0.4 ∆oct
1
ions, Oh field
ions, Oh field
eg
eg 1. What is the CFSE of [Fe(CN)6]3-?
eg
+ 0.6 ∆oct
+ 0.6 ∆oct
- 0.4 ∆oct
t2g
- 0.4 ∆oct
t2g
eg. 2. If the CFSE of [Co(H2O)6]2+ is -0.8 ∆oct,
what spin state is it in?
ions, Oh field
t2g
+ 0.6 ∆oct
Only
- 0.4 ∆oct
can be high or low spin
Jahn-Teller Distortion
configurations
Jahn-Teller Distortion
z
z-out: d-orbitals with z-component are stabilised
y
Oh field
x
+ 1/2 δ1
eg
- 1/2 δ1
∆oct
+ 2/3 δ2
t2g
The
complex
ligands contract or elongate giving a
- 1/3 δ2
2
Jahn-Teller Distortion
Jahn-Teller Theorum
z-in: d-orbitals with z-component are destabilised
"For a nonlinear molecule in an electronically degenerate state distortion must
occur to lower the symmetry, remove the degeneracy and lower the energy"
Oh field
x2-y2
+
eg
x2-y2 z2
1/
2
δ1
eg
x2-y2 z2
z2
- 1/2 δ1
t2g
yz
xz
xy
+ 1/3 δ2
yz
xz
Oh
z-in
Na2[CuF4] Oh
z-out
Cr2F5
z-out
xy
∆oct
t2g
e.g. K2[CuF4]
Oh
Cr(II) = high spin d4
yz xz
xy
- 2/3 δ2
e.g. [Cu(OH2)6]2+
Cu(II) = d9
z-out
Square planar geometry: ML4
Oh
z-out
square planar
x2-y2
eg
x2-y2
x2-y2 z2
z2
xy
t2g
xy
yz
xz
z2
xy
yz xz
yz xz
d8 complexes: some Ni(II), and Pd(II), Pt(II)
3
Tetrahedral crystal field splitting
Splitting of d-orbitals in a tetrahedral crystal field
eg
Δoct
metal ion
in free space
in symmetrical field
t2g
x2-y2 yz z2 xz xy
1
Effect of crystal fields on complex geometry
Spin orbit coupling is particularly significant for metals lower in the periodic
table but for 1st row transition metal complexes the formula can be used in a
simplified form because:
“The spin contribution outweighs the orbital angular momentum contribution to μeff"
Spin only magnetic moment, μSO
μSO =
n (n + 2)
where n = number of unpaired electrons
μSO = 2
S (S + 1)
where S = total spin quantum number = n / 2
2
Spin only magnetic moment, μSO
Measurement of magnetic susceptibility
μSO =
n (n + 2)
where n = number of unpaired electrons
Gouy Balance
What is the spin only magnetic moment of …..?
eg
x2-y2
z2
[Co(OH2)6]2+
n=
pivotal balance
μSO =
=
yz
xz
t2g
xy
yz
xz
t2
xy
[NiCl4]2μSO =
z2
=
x2-y2
[Ni(CN)4]2n=
sample
e
x2-y2
n=
scale
N
S
magnetic field
xy
μSO =
z2
=
yz
xz
3
Values of μeff for high spin Oh complexes
Effective magnetic moment, μeff
From experimentally measured molar magnetic susceptibility, χm
T is temperature in Kelvins
=
eh =
4 π me
9.27 x 10-24 J T-1
0
0
0
d1
1/
2
1.73
1.7 - 1.8
d2
1
2.83
2.8 - 3.1
d3
3/
2
3.87
3.7 - 3.9
Ti3+
Cr3+
Cr2+,
μB
d0
Sc3+, Ti4+
V2+,
2.83 is the magnetogyric ratio
μSO(BM)
dn
V3+
μeff in Bohr Magnetons:
S
M ion
μobs (BM)
Mn3+
d4
2
4.90
4.8 - 4.9
Mn2+, Fe3+
d5
5/
2
5.92
5.7 - 6.0
Fe2+, Co3+
d6
2
4.90
5.0 - 5.6
Co2+
d7
3/
2
3.87
4.3 - 5.2
Ni2+
d8
1
2.83
2.9 - 3.9
Cu2+
d9
1/
2
1.73
1.9 - 2.1
Zn2+
d10
0
0
0
4
Colour in TM complexes
Colour of d-d transitions depends on magnitude of Δ
[Ti(OH2)6]3+
Absorption spectrum: λmax = 510 nm
white light
400-800 nm
490-580 nm
The Colour Wheel
If red light is absorbed
Blue: 400-490 nm
yellow-green: 490-580 nm
Red: 580-800 nm
the complex appears green
If purple light is absorbed
the complex appears yellow
wavelength, λ (nm)
5
Effect of magnitude of Δ on colour
Effect of magnitude of Δ on colour
1. For a given ligand, the colour
2. For a given metal ion, the colour depends on
[V(H2O)6]3+
[V(H2O)6]2+
violet light absorbed
yellow light absorbed
complex appears yellow
complex appears violet
3+
2+
6
Colour and the Spectrochemical Series
Weak Field Ligands
Strong Field Ligands
High Spin Complexes
Low Spin Complexes
I- < Br- < S2- < SCN- < Cl-< NO3- < F- < OH- < ox2small Δ
< H2O < NCS- < CH3CN < NH3 < en < bpy
large Δ
< phen < NO2- < phosph < CN- < CO
7
Transition Metal Coordination Chemistry
[Ti(H2O)6]3+
Absorption spectrum: λmax = 510 nm
490-580 nm
Professor Sylvia Draper
eg
smdraper@tcd.ie
WHERE HAVE WE BEEN ?
Lecture 7: Magnetism and Colour
Tetrahedral crystal field
Magnetism
Spin only magnetic moment
Magnetism measurements
Absorption of light
d1 complexes and colour
Effect of ∆ on colour
WHERE ARE WE GOING ?
hν
t2g
eg
∆o
t2g
Lecture 8:
Colour continued, and Ligand Field
Theory
Selection rules
Limitations of CFT
LFSE and ∆
Factors affecting magnitude of ∆
1.
Oxidation state of metal e.g.
2.
Position of metal in periodic table
3.
Type of ligand in spectrochemical series
Thermodynamic aspects of LFT
1
Intensity of colour depends on …..
The Intensity of the colour depends on Selection Rules and
is measured by the absorbance
Absorbance is defined by the Beer-Lambert equation: A = log10(I0/I) = ε.c.l
where A = absorbance
I0 = intensity of incident light; I = intensity of light after passing through the cell;
ε = molar absorption coefficient; c = concentration; l = cell pathlength.
Typical Spectrum: Octahedral Cr(III) complex
Laporte Selection Rule
2.
Spin Selection Rule
At 590 nm, A = 0.138
A = εcl
0.138 = ε x 0.002M x 1 cm
ε = 0.138/ 0.002 x 1
ε = 69 M-1 cm-1
Absorbance
1.
Calculation of ε:
0.25
?
[Cr] = 0.002M
0.2
0.15
0.1
0.05
0
300 350 400 450 500 550 600 650 700
Wavelength (nm)
Transitions allowed by the selection rules (high ε values > 1000
M-1cm-1).
2
Laporte Selection Rule
Transitions between d-orbitals (g) forbidden by the Laporte selection rule
Octahedral complexes
Centrosymmetric: t2g and eg orbitals
Selection rule lifted by molecular vibration
s-orbital
gerade
e.g.
d-orbital
gerade
p-orbital
ungerade
p-orbital to d-orbital
by the Laporte selection rule
d-orbital to d-orbital
by the Laporte selection rule
Tetrahedral complexes
non-Centrosymmetric: t2 and e orbitals
Tetrahedral complexes are usually more strongly coloured than analogous
octahedral complexes
3
The Spin Selection Rule
Assumptions made in CFT
1.
The ligands are
2.
Interactions are purely
1.
Geometry
2.
Magnetism
3.
Colour
Explains
Limitations
Does not allow for
in M-L bonding
Doesn't explain
Order of ligands in the Spectrochemical Series
Nephelauxetic Effect
4
The Spectrochemical Series
The Nephelauxetic Effect
Weak Field Ligands
Strong Field Ligands
High Spin Complexes
Low Spin Complexes
"there is
in complexes than in the free ion"
- some
I- < Br- < S2- < SCN- < Cl-< NO3- < F- < OH- < ox2- < H2O < NCS- < CH3CN
< NH3 < en < bpy < phen < NO2- < phosph < CN- < CO
in M-L bonds – M and L share electrons
- effective size of metal orbitals
- electron-electron repulsion
Nephelauxetic series of ligands
F- < H2O < NH3 < en < [oxalate]2- < [NCS]- < Cl- < Br- < ICFT does not explain why some ligands with
charges are
than analogous ligands which are not charged
Nephelauxetic series of metal ions
Mn(II) < Ni(II) Co(II) < Mo(II) > Re (IV) < Fe(III) < Ir(III) < Co(III) < Mn(IV)
5
LFSE:the EXTRA stabilisation of complexes that occurs as a result of the splitting of orbitals in a ligand field
(cf. to CFSE: it is always positive and no account of P is taken)
LFSE Octahedral (always greater than tetrahedral) high/low spin options d4 to d7 only
eg
d1
d2
d3
d1 : LFSE 0.4
3/5∆o
-2/5∆O
t2g
d2 : LFSE 0.4 X 2 = 0.8
d3 : LFSE 0.4 x 3 = 1.2 maximum
d4 hs : LFSE 0.4 x 3 – 0.6 = 0.6
d4
d4
d5
eg
LFT is identical to CFT, except that it allows for some
in M-L bonding.
3/5∆o
-2/5∆O
t2g
parameters.
d5 hs : LFSE 0.4 x 3 – 0.6 x 2 = 0 minimum
d5 ls : LFSE 0.4 x 5 = 2.0 highest
d6 hs : LFSE 0.4 x 4 - 0.6 x 2 = 0.4
d6
LFT is based on
d4 ls : LFSE 0.4 x 4 = 1.6
d5
eg
t2g
d6
d7
d6 ls : LFSE 0.4 x 6 = 2.4
d7
3/5∆o
-2/5∆O
d7 hs : LFSE 0.4 x 5 – 0.6 x 2 = 0.8
d7 ls : LFSE 0.4 x 6 – 0.6 = 1.8
d8 : LFSE 0.4 x 6 – 0.6 x 2 = 1.2 maximum
ls gives higher LFSE than hs option
6
Variation in LFSE with ∆excluding P)
Lattice Energies (from Born Haber cycle)
For 1st row M(II) chlorides: hs, Oh
Ca(II)
ls d6:
LFSE
hs d0, d5, d10 :
LFSE
hs Oh d3, d8 :
LFSE
Ti(II)
V(II)
The energy
Cr(II)….etc
when ions come together from infinite
separation to form a crystal
Mn+(g) + n X-(g)
MXn(s)
Discrepancies in lattice energies are a result of
7
Variation in Hydration Enthalpies
Enthalpy of hydration:
M2+(g) + xs H2O
[M(OH2)6]2+(aq)
8
Variation in Ionic Radii
Ca2+
Zn2+
Sc3+
Ga3+
high spin
low spin
Last Lecture
9 lecture course on Coordination Chemistry
Prof. Sylvia Draper
high spin
low spin
So……. CFSE, LFE and MO theory agree !
- each one building upon and improving the other
low spin: steady
followed by increase from t2g6 eg1
high spin: steady
followed by increase from t2g3 eg1
1
Ionisation Energies of first row TM ions, hs, Oh
Irving-Williams series
[M(H2O)6]2+
+
[EDTA]4-
[M(EDTA)]2- +
6H2O
Position of equilibrium depends on
d5
d6
d7
d8
d9
d10
Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+
2
Valance Bond Theory
Tetrahedral e.g. [Zn(OH)4]2-
Square Planar e.g. [Ni(CN)4]2-
Ligand = Lewis base
Metal = Lewis acid
s, p and d orbitals give
with specific geometries
Number and type of M-L hybrid orbitals determines geometry of the complex
Octahedral Complex
e.g. [Cr(NH3)6]3+
Limitations of VB theory
Assumes bonding is
Does not account for
Can predict
wrongly
Does not account for
3
Molecular Orbital (MO) theory
Ligand Group Orbitals
MO theory considers covalent interactions
e.g. molecular hydrogen
Metal valance shell orbitals =
Linear Combination of Atomic Orbitals (LCAO)
Ligand valance shell orbitals =
LCAO approach
1s
e.g. [Co(NH3)6]3+ = sigma-bonded complex
1s
1s
zz
2s
2p
sp3 hybrid orbitals
N
H
H atom
H2 molecule
H atom
H
H
Six sp3 hybrid ligand orbitals form a set of
4
e.g. [CoF6]3-
Co3+ = 6 e-
t1u*
6 x F- = 12 ea1g*
Complexes with M-L π-bonding
e.g. H2O,OH-, lower halides Cl-, Br-, I-
Total =
high spin
electron density to the metal centre
ligand orbital and
4p
π-bond
σ-bond
π-bond
metal orbital
of electron density from filled p orbitals
4s
eg*
t2g
3d
6 x LGO
eg
π-acceptor ligands e.g. CO, N2, NO, alkenes
electron density from the metal centre
metal orbital and
t1u
O
C
C
O
ligand orbitals
L
Mn+
a1g
ML6n+
L
L
L
L
L
of electron density into emtpy π∗-antibonding orbitals
5
π-acceptor ligands: back-donation to empty ligand antibonding orbitals
t1u*
π-donor ligands: filled ligand orbitals and empty metal orbitals
t1u*
a1g*
small ∆oct
- weak field ligand
a1g*
Ligand π∗-orbitals
eg*
4p
large ∆oct
- strong field ligand
4p
Ligand π-orbitals
4s
(vacant)
eg*
4s
(occupied)
3d
3d
Ligand σ-orbitals
Ligand σ-orbitals
(occupied)
(occupied)
eg
Metal d electrons
(d6 example)
Metal d
electrons
t1u
(d6 example)
a1g
Energy
Mn+
ML6n+
eg
t1u
a1g
6L
Energy
Mn+
ML6n+
6L
6