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M.Sc. (Previous) Chemistry
Paper – I :
INORGANIC CHEMISTRY
BLOCK – I
UNIT – 1 : Stereochemistry in Main Group Compounds.
UNIT – 2 : Metal Ligand Bonding
UNIT – 3 : Electronic Spectra of Transition Metal
Complexes
Author – Dr. Purushottam B. Chakrawarti
Edtor – Dr. M.P. Agnihotri
1
UNIT-1 STEREO CHEMISTRY IN MAIN GROUP COMPOUNDS
Structure:
1.0
Introduction
1.1
Objectives
1.2
Valence Shell Electron Pair Repulsion (VSEPR)
1.3
1.4
1.5
1.6
1.2.1
Gillespie Laws
1.2.2
Applications
1.2.3
Comparison of CH4, NH3, H2O and H3O+
1.2.4
Comparison of PF5, SF4 and [ICl2-]
Walsh Diagrams
1.3.1
Application to triatomic molecules.
1.3.2
Application to penta atomic molecules.
d - p Bonds.
1.4.1
Phosphorous Group Elements.
1.4.2
Oxygen Group Elements.
Bent's rule and Energetic of hybridisation.
1.5.1
III and IV groups halides.
1.5.2
V and VI groups Hydrides halides.
1.5.3
Isovalent hybridisation.
1.5.4
Apicophilicity.
Some simple reactions of covalently bonded molecules.
1.6.1
Atomic invevsion.
1.6.2
Berrypseudo rotation.
1.6.3
Nucleophillic substitution.
1.7
Let Us Sum Up.
1.8
Check Your Progress: The key.
2
1.0
INTRODUCTION
The concept of molecular shape is of the greatest importance in
inorganic chemistry for not only does it affect the physical properties of
the molecule, but it provides hints about how some reactions might occur.
Pauling and Slater (1931) in their Valence Bond Theory (VBT) proposed
the use of hybrid orbital, by the central atom of the molecule, during bond
formation. Thus, with the knowledge of the hybridisation used by the
central atom of the molecule, one can predict the shape and also the
angles between the bonds of a molecule. However, since then more
advanced theories have come into existence. In this chapter we explore
some of the consequences of molecular shape in terms of VSEPR theory
and refine that concept into the powerful concept of molecular symmetry
and the language of group theory, using Walsh diagrams and Bent's
Theory, towards the end of the unit.
You may recall what you have already studied about the directional
property of a covalent bond and the concept of hybridisation of orbital to
predict molecular geometry.
1.1
OBJECTIVES
The main aim of this Unit is to study stereochemistry in main
group compounds. After going through this unit you should be able to:

discuss absolute shapes of various molecules;

explain anomalies in bond angles present in some molecules;

describe stability of molecular shapes in terms of energetics, and

discuss simple reactions of covalently bonded molecules on the
basis of the principles studied in this unit;
3
1.2
VSEPR THEORY (VALENCE SHELL ELECTRON PAIR
REPULSION THEORY)
In order to predict the geometry of covalent molecules, Valence
Shell Electron Pair Repulsion Theory is used. This theory was given by
Gillespie and Nyholm. According to this theory the geometry of a
molecule depends upon the number of bonding and non-bonding electron
pairs in the central atom. These arrange themselves in such a way that
there is a minimum repulsion between them so that the molecule has
minimum energy (i.e. maximum stability).
1.2.1 Gillespie Laws
The following rules have been reported by Gillespie to explain the
shape of some covalent molecules:
1.
If the central atom of a molecule is surrounded only by bonding
electron pairs and not by non-bonding electron pairs (lone pairs),
the geometry of the molecule will be regular.
In other words we can say that the shape of covalent molecule will
be linear for 2 bonding electron pairs, triangular for 3 bonding
electron pairs. tetrahedral for 4 bonding electron pairs, trigonal
bipyramidal for 5 bonding electron pairs:
Name of
Compound
BeCl2
Bonding
Electron Pairs
2
Shape
Linear
BeCl3
3
Triangular Planar
SnCl4
4
Regular Tetrahedral
PCl5
5
Trigonal bipyramidal
SF6
6
Regular Octahedral
4
2.
When the central atom in a molecule is surrounded by both,
bonding electron pairs as well as by lone pairs, then molecule will
not have a regular shape. The geometry of the molecule will be
disturbed. This alteration or distortion in shape is due to the
alteration in bond angles which arises due to the presence of lone
pairs on the central atom. How the presence of lone pairs causes an
alteration in bond angles can be explained as follows:
At a fixed angle the closer the electric-pairs to the central
atom, the greater is the repulsion between them. Since the lone-pair
electrons are under the influence of only one positive centre (i.e.
nucleus), they are expected to have a greater electron density than
the bond-pair electrons which are under the influence of two
positive centres. Thus lone pair is much closer to the central atom
than the bond pair. Hence it is believed that lone pair will exert
more repulsion on any adjacent electron pair than a bond pair will
do on the same adjacent electron pair.
(lp - lp) > (lp - bp)
....................(i)
(lp = lone pair and bp = bond pair)
If the adjacent electron pair is a bond pair, then repulsive
force between lone pair and bond pair will be greater than repulsive
force between two bond pairs.
(lp - bp) > (bp - bp)
....................(ii)
On combining relations (i) and (ii) we get
(lp - lp) > (lp - bp) > (bp - bp)
Thus the repulsion between two lone pairs is maximum in
magnitude, that between a bp and lp is intermediate while that
between two bond pairs is the minimum.
5
The more the numbers of lone pairs on a central metal atom,
the greater is the contraction caused in the angle between the
bonding pairs. This fact is clear when we compare the bond angles
in CH4, NH3 and H2O molecules. (Table 1.1)
Table 1.1
Molecules
3.
CH4
No. of Lone
pairs on
central atom
0
Bond
Angle
109.5o
Contraction in
bond angle
w.r.t. CH4
0
NH3
1
107.5o
2o
H2O
2
105.5o
4o
B-A-B bond angle decreases with the increase in electro negativity
of atom B in AB2 molecule where A is the central atom.
Example: Pl3 (102o) > P Br3 (101.5o) > PCl3 (100o)
4.
Bond angles involving multiple bonds are generally larger than
those involving only single bonds. However, the multiple bonds do
not affect the geometry of the molecule.
5.
Repulsion between electron pairs in filled shells are larger than the
repulsion between electron pairs in incompletely filled shells.
Examples: H2O (105.5o) < H2S (92.2o)
1.2.2 Applications of Gillespie Laws
Let us take some examples in support of these laws:
(a)
AX2 molecule, which has only two bond-pairs, will be linear:
X----A-----X
Examples in this groups will be BeCl2, CaCl2, CO2 etc.
6
(b)
If the molecule is AX3 (I) or AX2 with a lone pair of
electrons on the central atom A, i.e. AX2E (II), then the
molecule will be triangular (Fig 1.1):
Fig 1.1 :
(I) = BCl3, BF3 etc.
(II) = SO2, SnCl2 etc.
(c)
If the molecule is AX4 (III) or AX3E (IV) or AX2E2, then
AX4 will be tetrahedral; AX3E will be pyramidal and AX2E2
will be angular. (Fig. 1.2):
Fig 1.2 :
(III) = CCl4, CH4, SiCl4, GeCl4 etc.
(IV) = NH3, PCl3, As2O3 etc.
(V) = H2O, SeCl2, etc.
(d)
If the molecule is AX5 (VI) or AX4E (VII) or AX3E2 (VIII)
or AX2E3 (IX) then AX5 will be triangular bi pyramidal;
AX4E will irregular tetrahedral; AX3E2 will be T-shaped,;
and AX2E3 will be linear. (Fig. 1.3)
(VI)
(IX)
7
(VIII)
(IX)
Fig 1.3 : (VI) = PCl5; (VII) SF4, TeCl4 etc.
(VIII) = ClF3, BrF3 etc.
(IX) = XeF2, ICl 2 , or I 3 etc.
(e)
If the molecule is AX6 (X) or AX5E (XI) or AX4E2 (XII)
then AX6 will be octahedral, AX5E will be square pyramidal;
and AX4E2 will be square planar. (Fig. 1.4)
(X)
(XI)
Fig 1.4 :
(XII)
(X) =
SF6, WF6, etc.
(XI) =
BrF5, IF5 etc.
(XII) =
ICl4-, XeF 2 etc.
1.2.3 Comparison of CH4, NH3, H2O and H3O+
In table 1.1 bond angles in CH4, NH3 and H2O molecules are given.
In all these molecules, the central atom (C, N and O respectively) is sp3
hybridised. But they differ in the number of lone pair (s) present on the
central atom, being zero in CH4, one in NH3 and two in case of H2O.
Thus the repulsive force between electron pairs gradually increases in
these molecules from CH4 to H2O, resulting in the change of geometry
and the bond angles. This CH4 (Four bond pairs) is tetrahedral with the
8
characteristic bond angle of 109.5o. NH3 is pyramidal (Three bond pairs
and one lone pair) and has a bond angle of 107o. While H2O is angular
(Two bond-pairs and two lone-pairs) and has bond angle of 105o. The
increasing lp-lp repulsion decreases the bond angles from 109.5o to
~107o in NH3 and ~105.5o in H2O. On comparing H3O+ with these
molecules, we notice that, it resembles with NH3 molecule. As the central
atom of H3O+ on (Oxygen) is also sp3 hybridised and has 3bp+1 lp. Thus
H3O+ will also be pyramidal with a bond angle of ~107o. (Fig 1.5)
Fig 1.5 :A- CH4, Tetrahedral, bond angle 109.5o
B- NH4, Pyramidal, bond angle ~107o
C- H2O, Angular, bond angle ~105o
D- H3O+, Pyramidal, bond angle ~107o
1.2.4 Comparison of PF5, SF4, ClF3 and [ICl2-]
A comparison of PF5, SF4, ClF3 and [ICl2]- species clearly indicates
that each of these molecules have 10 electrons in the valence shell of the
central atom, being P, S, Cl and I respectively. In addition to this, the
central atom in each case is sp3d hybridised, and has 0, 1, 2 and 3 lone
pairs (lp) respectively. In PF5, all the five electron pairs (= 10 electrons)
are bond pairs and are housed in the five sp3d hybrid orbital; resulting in
the trigonal bipyramidal geometry of the molecule. (Fig 1.6 (a)).
9
SF4 molecule has 4 bond-pairs and one lone pair on the central S
atom. The lp in this molecule has two options- it can sit in a axial or in an
equatorial orbital. In the axial position (Fig. 1.6 (b))i)) it has three bps at
90o and one bp directly opposite to itself. While in equatorial position
(Fig. 1.6 (b)(ii)) it has two bps at 90o and two bps at 120o. As in the
equatorial position lp-bp repulsion is less and expansion is easy the lp
prefers the equatorial position and the molecule is therefore irregular
tetrahedral.
In ClF3., the two lps may be axial-axial (Fig 1.6(c)(i)), or axialequatorial (Fig 1.6(c)(ii)), or equatorial-equatorial (Fig 1.6(c)(iii))
positions. As the axial position will result in maximum repulsion hence
the axial position for the lp is ruled out. Thus the molecule will have T.
shaped geometry, according to the Fig. 1.6(c)(ii).
Similarly, due to reduced lp-lp repulsions and a larger volume that
a lp can occupy at equatorial positions, the [ICl2]- ion will be linear. The
three lps occupy the three equatorial positions, leaving the axial positions
for the Cl atoms (Fig. 1.6(d)).
(a) PF5, Trigonal bipyamidal
(I)
(II)
(b) SF4, Irregular tetrahedral
10
(c) ClF3, T-Shaped
(d) [ICl2]-, Linear
Fig. 1.6 : Structures of PF5, SF4, ClF3 and [ICl2]-
Limitations of VSEPR Theory
1.
This theory is not able to predict the shapes of certain transition
element complexes.
2.
This theory is unable to explain the shapes of certain molecules
with an inert pair of electrons.
3.
This theory is unable to explain the shapes of molecules having
extensive delocalised -electron system.
4.
This theory can not explain the shapes of molecules which have
highly polar bonds.
11
Check Your Progress-1
Notes: (i) Write your answers in the space given below;
(ii) Compare your answers with those given at the end of the unit.
(a)
Identify and write the shapes of the following molecules.
(i)
BeCl2.....................................(iv)
ClF3........................................
(ii)
PCl3.......................................(v)
XeF2........................................
(iii)
SF4.........................................(vi)
BrF5........................................ (b)
Arrange
the
following
species in the order of increasing bond-angle, considering the repulsive
forces due to lone pair:

NO2, NO2 and Θ NO2
>---------------->-------------->------------
1.3
WALSH DIAGRAMS
We know all the systems want to be in a stable state, and the stable
state is one in which it has the minimum possible energy. Same is true for
the stereochemistry or the geometry of a molecule. The VSEPR theory
considered that the most stable configuration of a molecule is one in
which repulsive forces between the valence electron-pairs is minimum. In
contrast, the molecular orbital theory (MOT) considers that the stable
geometry of a molecule can be determined on the basis of the energy of
molecular orbitals formed as a result of linear combination of atomic
orbitals (LCAO). In 1953 A.D.Walsh proposed a simple pictoralapproach to determine the geometry of a molecule considering and
calculating the energies of molecular orbitals of the molecule.
12
The basic approach is to calculate the energies of molecular
orbitals for two limiting structures, say linear or bent to 90 o for an AB2
molecule, and draw a diagram showing how the orbitals of one
configuration correlate with those of the other. Then depending on which
orbitals are occupied, one or the other structure can be seen to be
preferred. By means of approximate MO Theory implemented by digital
computers, this approach has been extended and generalized in recent
years.
Walsh's approach to the discussion of the shape of an AB2
triatomic molecule (such as BeH2 and H2O) is illustrated in Fig. 1.8. The
illustration shows an example of a Walsh diagram, a graph of the
dependence of orbital energy on molecular geometry. A Walsh diagram
for an B2A or AB2 molecule is constructed by considering how the
composition and energy of each molecular orbital changes as the bond
angle changes from 90o to 180o. The diagram is in fact just a more
elaborate version of the correlation diagram.
1.3.1 Application to Triatomic Molecules
The coordinate system for the AB2 molecule is shown in Figure
1.8. The AB2 molecule has C2v symmetry when it is bent and, when linear
D2h symmetry. To simplify notations, however, the linear configuration is
considered to be simply an extremum of the C2v symmetry. Therefore the
labels given to the orbitals through the range 90 o ≤ θ < 180o are retained
even when θ = 180o. The symbols used to label the orbitals are derived
from the orbital symmetry properties in a systematic way, but a detailed
explanation is not given here. For present purposes, these designations
may be treated simply as labels. (Fig. 1.7).
13
Fig. 1.7
The A atom of AB2 molecule will be assumed to have only s, px,
py and pz orbitals in its valence shell, whereas each of the B atoms is
allowed only a single orbital oriented to form a σ bond to A. In the linear
configuration PAX and PAZ are equivalent non-bonding orbitals labelled
2a, and b1 respectively. The orbitals SA and PAY interact with σ1B and σ2B,
σ orbitals on the B atoms, to form one very strongly bonding orbital, 1a,
one less strongly bonding orbital, 1b2, one less strongly bonding 3a1 and
3b2. The ordering of these orbitals and in more detail, the approximate
values of their energies can be estimated by an MO calculation. Similarly,
for the bent molecule the MO energies may be estimated. Here only p zA is
non bonding, spacing and even the order of the other orbitals is function
of the angle of bending θ. The complete pattern of orbital energies, over a
range of θ, is obtained with typical input parameters. This is shown in the
figure 1.8. Calculations in the Huckel approximation are simple to
perform and give the correct general features of the diagram but for
certain cases (e.g. AB2E2) very exact computations are needed for an
unambiguous prediction of structure.
14
(Angle θ)
Figure 1.8 Orbital Correlation Diagram For AB2 Triatomic
Molecules Where A uses only s and p orbitals
Form the approximate diagram (Fig. 1.8) it is seen that an AB2
molecule (one with no lone pairs) is more stable when linear then when
bent. The 1b2 orbital drops steadily in energy form θ = 90o to 180o; while
the energy of the 1a1 orbitals is fairly insensitive to angle.
For an AB2E molecule the results are ambiguous, because the trend
in the energy of the 2a1 orbital approximately offsets that of the 1b2
orbiral.
For AB2E2 molecules, the result should be the same as for AB2E.
Since the energy of b1 orbirtal is independent of the angle. Thus it is not
clear in this approach that AB2E2 molecules should necessarily be bent,
but all known ones are.
15
The H2O Molecule:
Because of its unique importance, this molecule has been subjected
to more detailed study than any other AB2E2 molecule. A correlation
diagram calculated specially for H2O is shown in the figure. Although it
differs in detail for the general AB2E2 shown in the figure it is
encouraging to see that the important qualitative features are the same.
The general purpose diagram pertains to a situation in which there is only
a small energy difference between the ns and np orbitals of the central
atom. As stated in discussing that general purpose diagram, it is not clear
whether an AB2E2 molecule ought necessarily to be bent.
In the diagram calculated expressly for H2O the lowest level's is
practically pure 2s and its energy is essentially constant for all angles. It
can be determined from this diagram that the energy is minimized at an
angle of 106o, essentially in accord with the experimental value of 104.5o.
(Fig. 1.9).
Fig. 1.9
16
BeH2 - Molecule
The simplest AB2 molecule in Period 2 is the transient gas-phase
BeH2 molecule (BeH2 normally exists as a polymeric solid), in which
there are four valence electrons. These four electrons occupy the lowest
two molecular orbitals. If the lowest energy is achieved with the molecule
angular, then that will be its shape. We can decide whether the molecule
is likely to be angular by accommodating the electrons in the lowest two
orbitals corresponding to an arbitrary bond angle in Fig. 1.7. We then
note that the HOMO decreases in energy on going to the right of the
diagram and that the lowest total energy is obtained when the molecules
is linear. Hence, BeH2 is predicted to be linear and to have configuration
1 g2 2 u2 . In CH2, which has two more electrons than BEH2, three of the
molecular orbital must be occupied. In this case, the lowest energy is
2
2
2
achieved if the molecule is angular and has configuration 1 1a1 2a111b2 .
The principal feature that determines whether or not the molecule
is angular is whether the 2a1 orbital is occupied. This is the orbital that
has considerable A-2s character in the angular molecule but not in the
linear molecule.
1.3.2 Application to Penta Atomic Molecules
For peta-atomic molecules examples of CH4 and SF4 may be taken
for consideration. For these molecules, two geometries are possible: one a
symmetrical tetrahedral and the other, a distorted tetrahedral geometry (or
a tetragonal geometry of relatively lower symmetry).
CH4 - Molecule
Methane, CH4, has eight valence electrons. During bonding, for
orbitals [a1g (2s) and t1u (2p)] of carbon, and one [a1g (1s)] orbital of each
17
four hydrogen atoms take part. Overlapping of these eight orbitals, eight
molecular orbitals are formed, the four bonding (2σg or 2a 1 and 2tσ or
2t1) and the four antibonding (2σg* and 2t*σ). The eight valency
electrons of CH4 molecule are distributed in the four bonding molecular
orbitals (Electronic Configuration, 2 12 2t1b ). In the tetrahedral geometry
due to overlapping with the orbitals of hydrogen atom the energy of 2a1g
and 2t1u orbitals is considerably reduced. In contrast, in the distorted
geometry, comparatively less overlapping of t 1u orbitals with the
hydrogen orbitals (as compared to that in the tetrahedral geometry) the
energy of 2t1 molecular orbitals increases. Thus the geometry of CH4
molecule is symmetrical tetrahedral, rather than a distorted tetrahedral.
SF4 - Molecule
In the valence shall of sulphur atom, in SF4 molecule, in addition to
3s (α1g) and 3p (t1u) orbitals, 3d (t2g and eg) orbitals are also present.
During the bonding, 2pz orbital of each of the four fluoring atoms take
part. As a result for bonding (2a1 and 2t1) and four antibonding (2 1* 2t1* )
molecular orbitals formed; while the d-orbitals [α1(dz2), b1(dx2 - y2) and t2
(dxy, dxz, dyz)] are present as non-bonding molecular orbitals. Ten valency
electrons of SF4 molecule remain distributed in the four bonding and one
non-bonding molecular orbitals, resulting in 2 12 ,2t1b . 3a12 configuration.
As in the distorted geometry, overlapping of 2t 1 orbitals is comparatively
greater (thus reducing their energy) and the filling in 3α1 orbital
considerably reduce the energy of the system as compared to that in the
regular tetrahedral structure. Hence SF4 molecule has a distorted
tetrahedral geometry, rather than a regular tetrahedral structure.
Thus we can say 'Walsh Diagrams' are complementary to the
VSEPR concept.
18
Check Your Progress 2
Notes:1.
2.
Write your answers in the space given below.
Compare your answers with those given at the end of this
unit.
(a).
Predict the shape of an H2O molecule on the basis of a
Walsh Diagram for an AB2 molecule.
-----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
(b)
Is any AB2 molecule, in which A denotes an atom of a
period 3 element, expected to be linear? If so, which
-----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
1.4
d - p BONDS.
There are several structural phenomena that have traditionally been
attributed to the formation of d - p Bonds. Recent work has raised
some doubts. The phenomena in questions are exemplified by:
1.
The fact that for amines such as (R3Si)2 NCH3, (R3Si)3 N and
(H3Ge)3 N, the central NSi2C, NSi3 and NGe3 skeletons are planar.
2.
Many tetrahedral species such as SiO4-4, PO4-3, SO4-2 and ClO-4
have bond lengths shorter than those predicted from conventional
tables of single bond radii. In silicates the Si-O-Si units also show
what were considered to be Si-O distance that are "too short" for
single bonds.
19
Recent re-examinations of these phenomena by both theoretical
and experimental methods together with earlier arguments now suggest
that the d - p contributions to these effects are at best small. Thus, in
reading literature written prior to 1985, where such interactions are often
accorded great importance, one should now be sceptical of all but the
facts themselves.
This is not to say that d - p bonding in main group compounds is
never important. Probably in the case of - S - N = S units, and in F3F ≡ N,
where the S - N distances are very short indeed. However, it is always
dangerous to attribute all structural effects to simple orbital overlaps,
even if the explanation seems to fit, and the rise and fall of the d - p
overlap hypothesis is a case in point.
In a multiple bonded molecule having bond pairs + lone pairs = 4,
5 or 6, the  bonds will be p - d bond. In this case the central atom
uses all its p-orbitals for hybridisations and has only d-orbitals available
to overlap with p-orbitals of the adjacent atom to give d- p bond.
The formation of d - p bond is common for all the second period
elements and is not important for the elements of third and higher
periods. The p-d bonding is more favourable than the d - p bonding
for higher atoms of third and higher periods.
1.4.1 d - p Bonding in Phosphorous Group Elements
Phosphine Oxide, R3P = 0, presents an important example of the
participation of d-atomic orbitals of nonmetallic elements in  bonding.
Presence of  bonding is defected with the help of the evidences, such as
reduction in the bond length, increase in the bond strength and the
stabilisation of charge distribution. On these grounds, compared to
20
ammine oxide, phosphine oxide presents a strong evidence of the
presence of d - p bond, in a very high stability of P = O.
Similarly the fact that almost all t-phosphines are readily oxidised
into R3P = 0, also indicates that d- p bond is present in the P = O
bonding. This is supported by the lower dipole moment of triethyl
phosphine oxide (1.4 x 10-3 Cm. cf 16.7 x 10-3 Cm. of trimethyl amine
oxide), higher dissociation energy of P = O bond (500-600 KJ, cf 200300 KJ of NO bond) and the smaller P-O bond lengths in phosphoryl
compounds.
1.4.2 d-p Bonding in Nitrogen, Oxygen and Sulphur Compounds
There are number of examples, which show d- p bonding in
nitrogen, oxygen and sulphur compounds:
(i)
Mobile  bonding in trisilyl amine results in the resonance in the
molecule :
_ +
H3Si = N
SiH3
SiH3
(ii)
The
bond
angles
in
disiloxane,
H3SiOSiH3,
and
silyl
isothiocyanate, H3Si-N=C=S indicates p - d back-bonding (in
comparison to ether and methyl isothiocyanate) :
Dimethyl ether
(Sp3-hybridised. 0 + 2 lp)
Addition compound with BF3
21
Disiloxane
(Sp3-hybridised. 0 + 1 lp)
No addition compound withBF3
Methyl isothicyanate
(one lp on N; Angular)
Silyl isothiocynate
lp used in  bonding, Linear
The argument given against the use of d atomic orbitals in bonding
by non-metallic elements is that a very high excitation energy is required
for the same. Hence it may be concluded that the use of d-atomic orbitals
in bonding by non-metallic elements will be possible only in their higher
oxidation states and when they are linked with strong electronegative
elements, eg. PF5,SF6,OPX3 etc.:
(i)
The N - S bond length in N  SF3 indicates. The bond order = 2.7
indicating d - p bonding:
Thiazytrifluoride
N - S bond length = 141.6 pm; Bond order = 2.7
(cf N - S = 174 pm, b.0 = 1, N = S = 154 pm; b.0 = 2)
(ii)
In S4N4F4 also there is indication of d‫ ה‬- p‫ ה‬bonding (compare with
S4H4N4):
Tetra Sulphur tetramide
(Isoelectronic and
Isomorphous with S8
Tetra sulphur tetramide fluoride
(Alternate S = N bond)
22
(ii)
In diphenyl phosphonitrilic fluoride, there is evidence of ‫ ה‬bonding
in the ring:
Check Your Progress-3
Notes: 1.
2.
Write your answers in the space given below.
Compare your answers with those given at the end of this
unit.
(a).
In which of the following molecule there is a possibility of
d-p bonding?
S4N4F4, S4H4N4, N ≅ SF3, H3SiOSiH3, CH3N = C=S and
Et3NO.
Ans: (i) -------------------------------
(ii)-----------------------
(iii) ----------------------------(b)
Which are the evidences in favour of d-p bonding
in R3PO molecule?
Ans: (i) ---------------------------------(ii) --------------------------------(iii) --------------------------------
23
1.5
BENT'S RULE AND ENERGETICS OF HYBRIDISATION:
BENT RULE
Bent rule may be stated as follows:
"More electronegative constituents 'prefer' hybrid orbitals having
less s character and more electropositive substituents 'prefer' hybrid
orbitals having more s character."
An example of Bent rule is provided by the fluoromethanes. In
CH2F2, the F-C-F bond angle is less than 109.5o, indicating less than 25%
s character, but the H-C-H bond angle is larger and C-H bond has more s
character. The bond angle in the other fluoromethanes yield similar
results.
The tendency of more electronegative substituents to seek out the
low electronegative pxdx2 apical orbital in TBP structures is often termed
"apicophilicity". It is well illustrated in a series of oxysulfuranes of the
type-
prepared by Martin and Co-workers. These, as well as related
phosphoranes provide interesting insight into certain molecular
rearrangements.
Bent's rule is also consistent with and may provide alternative
rationalization for Gillespie's VSEPR model. Thus the Bent's rule
prediction that highly electronegative constituents will 'attract' p character
and reduce bond angles is compatible with the reduction in regular
24
volume of the bonding pair when held tightly by an electronegative
substituents. Strong, s-rich covalent bonds require a larger volume in
space to bond. Thus double bonded oxygen, despite the high electro
negativity of oxygen, seeks s-rich orbitals because of the shortness and
better overlap of the double bond. Again, the explanation, whether in
purely s-character terms (bent's rule) or in larger angular volume for a
double bond (VSEPR), predicts the correct structure.
The mechanism operating behind Bent's rule is not completely
clear. One factor favouring increased p character in electronegative
substituents is the decreased bond angles of p orbitals and the decreased
steric requirements of electronegative substituents. There may also be an
optimum strategy of bonding for a molecule in which the character is
concentrated in those bond in which the electronegativity difference is
small and covalent bonding is important. The p character, if any, is then
directed towards bonds to electronegative groups. The latter will result in
greater ionic bonding in a situation in which covalent bonding would be
low anyway because of electronegativity difference.
Some light may be thrown on the workings of Bent's rule by
observations of apparent exceptions to it. The rate exceptions to broadly
useful rules are unfortunate with respect to the universal applications of
those rules. They also have the annoying tendency to be confusing to
someone who is encountering the rule for the first time. On the other
hand, any such exception or apparent exception is a boon to the research
since it almost always provides insight into the mechanism operating
behind the rule.
25
Consider the cyclic bromophosphate ester.
The phosphorus atom is in an approximately tetrahedral
environment using four σ bonds of approximately sp3 character. We
should expect the more electronegative oxygen atoms to bond to s-poor
orbitals on the phosphorus and the two oxygen atoms in the ring do
attract hybridizations of about 20%s. The most electropositive constituent
on the phosphorus is the bromine atom and Bent's rule would predict an
s-rich orbital, but instead it draws another s-poor orbital on the
phosphorus atom is that involved in σ bond to the exocyclic oxygen. This
orbital has nearly 40% s-character. The oxygen atom ought to be about as
electronegative as the other two, so why the difference? The answer
probably lies in the overlap aspect.
1.
The large bromine atom has diffuse orbitals that overlap poorly
with the relatively small phosphorus atom. Thus, even though the
bromine is less electronegative than the oxygen, it probably does
not form as strong a covalent bond.
2.
The presence of a  bond shortens the exocyclic double bond and
increases the overlap of the σ orbitals. If molecules respond to
increase in overlap by rehybridization in order to profit from it, the
increased s-character then becomes reasonable. From this point of
view, Bent's rule might be rewarded. The p character tends to
concentrate in orbitals with weak covalently (from either electro
negativity or overlap considerations), and s-character tends to
26
concentrate in orbitals with strong covalently matched electro
negativities and good overlap.
Some quantitative support for the above qualitative arguments
comes from average bond energies of phosphours, bromine and oxygen.
P-Br
264 KJ Mol-1
P-0
335 KJ Mol-1
P=0
544 KJ Mol-1
Bent's rule is a useful tool in inorganic and organic chemistry. For
example, it has been used to supplement the VSPER interpretation of the
structures of various non-metal fluorides, and should be applicable to a
wide range of question on molecular structure.
Energetics of Hybridisation:
According to hybridisation model, bond directions are determined
by a set of hybrid orbitals on the central atom which are used to form
bonds to the ligand atoms and to hold unshared pairs. Thus AB2
molecules are linear due to the use of linear sp hybrid orbitals. AB 2
molecule should be equilaterally triangular, while AB2E molecule should
be angular, due to use of trigonal sp2 hybrids. AB4, AB3E and AB2E2
molecule should be tetrahedral, pyramidal and angular, respectively,
because here sp3 hybrid orbitals are used.
These cases are, of course, very familiar and involve no more than
an octet of electrons.
For the AB5, AB4E, AB3E2 and AB2E3 molecules the hybrid must
now include orbitals in their formation. The hybrid orbitals used must be
of the sp3dz2 leading to TBP geometry and Sp3dx2-y2 leading to SP
geometry. There is no way to predict with certainty which set is preferred,
and doubtless that difference between them connot be great, since we
27
know experimentally that AB5 molecules nearly all have TBP structures,
the same arrangement is assumed for the AB4E cases, and so on. Even
this adhoc assumption does not solve all difficulties, since the position
preferred by lone pairs must be decided and there is no simple physical
model here (as there was in the VSEPR approach) to guide us. A
preference by lone pairs for equatorial positions has to be assumed. With
these assumptions, a consistent correlation of all structures in this fiveelectron-pair class is possible.
For AB6 molecule, octahedral sp3d2 hybrids are used. For AB4E2
molecule; there is nothing in the directed valence theory itself to show
whether the lone pairs should be cis or trans. The assumption that they
must be trans leads to consistent results.
The most fundamental problem with the hybridisation model is that
in all cases in which there are more than four electron pairs in the valence
shell of the central atom, it is necessary to postulate that at least one d
orbital becomes fully involved in the bonding. There are both
experimental and theoretical reasons for believing that this is too drastic
an assumpiton. Some recent MO calculations and other theoretical
considerations suggest that although the valence shell d orbitals make a
significant contribution to the bonding in many cases, they never play as
full a part as do the valence shell p orbitals. Fairly directed experimental
evidence in the form of nuclear quadruple resonance studies of the I Cl 2
and ICl 4 3 ions shows that in these species, d-orbitals participation is very
small. This participations is probably greater in species with more
electronegative ligand atoms such as PF5, SF6 and Te (OH)6 but not of
equal importance with the contribution of the s and p orbitals.
28
Perhaps it is surprising that by going to the opposite extreme,
namely by omitting all consideration of d orbitals, but still adhering to the
concept of directed orbitals it is again possible to rationalize many of the
principal features of the structures of main group.
Ʃ Es+p3 = [2(-1806)]+[3(-981)] = 6555 KJ mol-1
For tetrahedral hybridised phosphours.
(3 te22te13te1) the energy will be:
Ʃ te = 5x (-1187) = -5935 KJ mol-1
In this case the hybridisation has cast 620 KJ mol -1 of energy or
roughly two bonds worth of energy. This is shown graphically in the
figure 1.10.
Figure 1.10
The energy difference between the hybridised and unhybridised
atom represents the increase in energy of the two electrons in the filled 3s
orbital and the decrease in energy of the electrons in the half-filled 3p
orbitals.
The energetics of hybridisation, together with the principle of good
overlap, are important in determining the electronic structure of
molecules.
29
1.5.1 Third and Fourth Groups Halides
For promoting an atom to hybridised excited state energy is
required. But when the hybrid orbitals give very strong bonds, the
energy gained from bonding may be used for excitation. For
example, when carbon forms four covalent bonds, although there is
a promotion energy from Is22s22p2  ls22s12p3, this is independent
of the hybridization to the valence state :
Figure 1.11 Hybridisation energy of carbon
This energy of hybridisation is the order of bonding energy. Its
important use is to determine structure of molecules. For example, the
stability of the halides of the III A (Gr. 13) and IV A (Gr. 14) viz BCl3,
AlCl3, GaCl3, InCl3, TiCl3 and CCl4, SiCl4, GeCl4, SnCl2, PbCl2, can be
explained on the basis of hybridisation energy. The heaviest elements of
these groups (Thallium and, Tin and Lead) the stable oxidation-states are
two unit less than the maximum oxidation states (3 and 4 respectively)
i.e. One in Thallium, TlCl and two in Tin and Lead, SnCl 2 and PbCl2.
Although, in these elements (Compared to the light elements of the
group) excitation is easy. It is because, in the heavier elements orbitals
are more diffused (Fig 1.12), hence in a volume region, values of  Tl or
 Pb are less, compared to  B or  Si. This results in lesser overlapping
of the central atom orbitals with orbitals of chlorine atoms. Hence, in the
compounds of heavy elements bonds are weaker, compared to that of the
30
light elements. Because of this less effective overlapping, the heavy
elements utilise pure p-orbitals in bonding and letting the lone-pair 'sink'
into a pure s-orbital. Hence, the stable chlorides of thallium and lead are
TlCl and PbCl2.
Figure 1.12 Effect of orbital size on overlapping
1.5.2 Fifth and Sixth groups Hydrides & Halides
As the energy of hybridization is of the order of magnitude of bond
energies and can thus be important in determining the structure of
molecules. It is responsible for the tendency of some lone pairs to occupy
spherical,
nonstereochemically
active
s-orbitals
rather
than
stereochemically active hybrid orbitals. For example, the hydrides of the
Group VA (15) and VIA (16) elements are found to have bond angles
considerably reduced as one progresses fro m the first element in each
group to those that follow (table 1.2 ). An energy factor
Table 1.2: Bond angles in the hydrides of Groups VA (15) and VIA (16)
NH3 = 107.2o
PH3 = 93.8o
AsH3 = 91.8o
SbH3 = 91.3o
OH2 = 104.5o
SH2= 92o
SeH2 = 91o
TeH2 = 89.5o
that favours reduction in bond angle in these compounds is the
hybridization discussed above. It costs about 600 KJ mol -1 to hybridize
31
the central phosphorus atom. From the standpoint of this energy factor
alone the most stable arrangement would be utilizing pure p-orbital in
bonding and letting the lone pair "sink" into a pure s-orbital. Opposing
this tendency is the repulsion of electrons, both bonding and nonbonding
(VSEPR). This favours an approximately tetrahedral arrangement. In the
case of the elements N and O the steric effects are most pronounced
because of the small size of atoms of these elements. In the larger atoms,
such as those of P, As, Sb, S, Se and Te, these effects are somewhat
relaxed, allowing the reduced hybridization energy of more p character in
the bonding orbitals to come into play. The molecule is thus forced to
choose between higher promotion energies and better overlap for an-srich hybrid, or lower promotion energies and poorer overlap for an s-poor
hybrid. (s-character: sp3 (25%)<sp2(-33%)<sp(50%)); s-rich means>25%;
s-poor means<25%)
1.5.3 Isovalent Hybridisation
In many tetravalent molecules the bond angle is seen slightly
distorted, than the ideal 109o28' for example in CH3Cl, H-C-H bond is
110o20'. This deviation may be explained in terms of 'Isovalent
Hybridisation". Consider an imaginary molecule A-M-B. If in this
molecule, B is replaced by a strong electronegative element C, then M is
rehybridised in such a way that the orbital used for bonding C has more p
character, than the orbital used for bonding B. Hence in A-M-C molecule,
M-A bond will have more s-character, compared to that in A-M-B
molecule. For example, as compared to CH3NH2, the carbon atom in
CH3OH uses a hybrid orbital having more s-character to link methyl
hydrogen. As a result the C-O bond has more p character as compared to
the C-N bond. This is an example of Bent's rule.
32
1.5.4 Apicophilicity
Good example of the effect of the differences in hybrid bond
strengths are shown by the bond lengths in MXn Molecules with both
equatorial and axial constituents. (Table 1.3)
Table 1.3
req(pm)
rax(pm)
PF5
153.4
157.7
PCl5
202
214
SbCl5
231
243
SF4
154
164
ClF3
159.8
169.8
BrF3
172.1
181.0
An sp3d hybrid orbital set may be considered to be a combination
of pzdz2 hybrids and spxpy hybrids. The former make two linear hybrid
orbitals bonding axially and the latter form the trigonal, equatorial bonds.
The sp2 hybrid orbitals are capable of forming stronger bonds, and they
are shorter than the weaker axial bonds. When the electronegativities of
the substituents on the phosphorus atom differ, as in the mixed
chlorofluorieds, PClxF5-x, and the alkylphosphorus fluorides, RxPF5-x it is
experimentally observed that the more electronegative substituent
occupies the axial position and the less electronegative substituent is
equatorially situated. This is an example of Bent's rule which states: More
electronegative substituents 'prefer' hybrid orbitals having less scharacter, and more electropositive substituents 'prefer' hybrid orbitals
having more s-character.
A second example of Bent's rule discussed earlier is that of the
fluoromethanes. In CH2F2 the F-C-F bond angle is less than 109.5o,
33
indicating less than 5% s character, but the H-C-H bond angle is larger
and the C-H bond has more s character. The bond angles in the other
fluoromethanes yield similar results.
The tendency of more electronegative substituents to seek out the
low electronegativity pzdz2 apical orbital in TBP structure is often termed
"apicophilicity". It is well illustrated in a series of oxysulfuranes.
Check Your Progress-4
Notes: 1.
2.
Write your answers in the space given below.
Compare your answers with those given at the end of this
unit.
(a).
Bent's rule is exemplified in(i) Monovalency of.................and divalency of.......
(ii)
Decreasing bond angles of ................................
(iii)
F-C-F bond angle is .......................than 109o28'
but H-C-H bond angle is ........................ than
109028' in CH2F2 molecule.
(b)
Apicophiliaity is .........................................................
.....................................................................................
It is well illustrated by...............................................
1.6
SOME SIMPLE REACTIONS OF COVALENTLY BONDED
MOLECULES
One of the major differences between organic and inorganic
chemistry is the relative emphasis placed on structure and reactivity.
Structural organic chemistry is relatively simple, as it is based on
diagonal, trigonal or tetrahedral
carbon. Thus organic chemistry has
turned to the various mechanisms of reaction as one of the more exciting
34
aspects of the subject, to contrast, inorganic chemistry has a wide variety
of structural types to consider, and even for a given element there are
many factors to consider. Inorganic chemistry has been, and to a large
extent still is more concerned with the static structure of reactants or
products than with the way in which they interconvert. This has also been
largely a result of the paucity of unambiguous data on reaction
mechanisms. However, this situation is changing. Interest is increasingly
centring on how inorganic molecules change and react. Most of this work
has been done on coordinate chemistry, and much of it will be considered
later on, but a few simple reactions of covalent molecules will be
discussed here.
1.6.1 Atomic Inversion
The simplest reaction is seen in a molecule of ammonia. This can
undergo the simple inversion of the hydrogen atoms about the nitrogen
atom. This is analogous to the inversion of an umbrella in a high wind.
One might argue that above equation does not represent a reaction
because the product is identical to the reactant and no bonds were formed
or broken in the process. Leaving aside, the process illustrated above is of
chemical interest and worthy of chemical study.
Consider the trisubstituted amines and phophines shown in the
figure below.
(Chiral amines and phophines)
35
Because these molecules are non superimposable upon their mirror
images (i.e. they are chiral) they are potentially optically active, and
separation of the enantiomers is at least theoretically possible.
Racemization of the optically active material can take place as shown in
mechanism of NH3. It is of interest to note that the energy barrier to
inversion is strongly dependent on the nature of the central atom and that
of subsequent. For example, the barrier to inversion of methyl propyl
phenylphosphine is about 120 MJ Mol-1. This is sufficient to allow the
separation of optical isomers, and their racemization may be followed by
classical techniques. In contrast, the barrier to inversion in most amines is
low (-40 KJ mol) with such low barriers to inversion, optical isomers
cannot be separated because racemization takes place faster than the
resolution can be affected. Since traditional chemical separations cannot
effect the resolution of the racemic mixture, the chemist must turn to
spectroscopy to study the rate of interconversion of the enantiomers.
1.6.2 Berry Pseudo Rotation
In PF5 the fluorine atoms are indistinguishable by means of NMR
of F. This means that they are exchanging with each other faster than the
NMR instrument can distinguish them. The mechanism for this exchange
is related to the inversion reaction we have seen for amines and
phosphines. The mechanism for this exchange is believed to take place
through conversion of the ground state trigonal bipyramidal into a square
pyramidal transition state and back to a new trigonal bipyramidal
structure. This process results in complete scrambling of the fluorine
atoms at the equatorial and axial positions in phosphorus pentafluoride. If
it occurs faster than the time scale of NMR experiment, all the fluorine
atoms appear to be identical. Because it was first suggested by Berry, and
because, if all of the substituents are the same as in PF5, the two triogonal
36
bipyramidal arrangements are related to each other by simple rotation, the
entire process is called a Berry pseudorotation. Note that the process
can take place very readily because of the similarity in energy between
trigonal bypyramidal and square pyramidal structures.
(Berry pseudorotation in Pentavalent Phosphorus Compound)
In fact the series of 5-coordinated structures collected by
Muetterties and Guggenberger, which are intermediate between trigonal
bipyramidal and square pyramidal geometrically effectively provides a
reaction coordinate between the extreme structures in the Berry
pseudorotation.
1.6.3 Nucleophillic Substitution
The simplest reaction path for neucleophilic displacement may be
illustrated by solvolysis of a chlorodialkylphophine oxide.
We would expect the reaction to proceed with inversion of
configuration of the phosphorus atom. This is generally observed
especially when
the
entering
and
leaving
groups
are
highly
electronegative and is thus favorably disposed at the axial positions, and
when the leaving group is one that is easily displaced. In contrast in some
cases when the leaving group is a poor one, it appears as though front
37
side attack takes place because there is retention of configuration. In
either case, the common inversion or the less common retention, there is a
contrast with the loss of stereochemistry associated with a carbonium ion
mechanism.
1.6.4 Free Radical Mechanism
In the atmosphere there are many free radical reactions initiated by
sunlight. One of the most important and controversial sets of atmospheric
reactions at present is that revolving around stratospheric ozone. The
important of ozone and the effect of ultraviolet radiation on life will be
discussed later, but we may note briefly that only a small portion of the
sun's spectrum reaches the surface of the earth and that parts of UV
portion that are largely screened can cause various ill effects to living
systems.
The earth is screened from extremely high energy UV radiation
cleaves the oxygen molecule to form two free radicals of oxygen atoms.
O2 + hv (below 242 mm)  O. + O.
The oxygen atoms can then attack oxygen molecules to form
ozone.
O. + O2 + m  m + O3
The neutral body m carries off some of the kinetic energy of the
oxygen atoms. This reduces the energy of the system and allows the bond
to form to make ozone. The net reaction is therefore:
3O2 + hv  2O3
The process protects the earth from the very energetic, short
wavelength UV radiation and at the same time produces ozone, which
absorbs somewhat longer wavelength radiation by similar process:
O3 + hv (220-230 mm)  O2 + O.
Thus the process is repeated.
38
Check Your Progress-5
Notes: 1. Write your answers in the space given below.
2. Compare your answers with those given at the end of this unit.
(a).
During atomic inversion of amines, separation of optically active
isomers is not possible: because ...........................................................................................................
............................................................................................................
(b).
During Berry Pseudorotation process in PF5, the mechanism is
believed to take place through.
............................................................................................................
............................................................................................................
(c).
Solvolysis (Alcoholysis) of chlorodialkyl phosphineoxide proceeds
with inversion of configurations of phosphorous atom, when the
entering and leaving groups are ............................................................................................................
............................................................................................................
1.7
LET US SUM UP
After going through this unit, you would have achieved the
objectives stated earlier in this unit. Let us recall what we have discussed
so for:

VSEPR theory considers that the geometry of a molecule depends
upon the number of bonding and nonbonding electron pairs in the
central atom. These arrange themselves in such a way that there is
a minimum repulsion between them so that the molecule has
minimum energy (i.e. maximum stability).
39
The repulsive force between the electron pairs very as:
(lp - lp) > (lp - bp) > (bp - bp)
Thus, the more that number of lone pair on a central metal atom,
the greater is the contraction caused in the angle between the
bonding pair. Hence, the bond-angle in CH4 (4bp + Olp) is 109o28',
in NH3 (3bp + lp) is 107o and H2O (2bp + 2lp) is 105o. Similarly.
BCl3 is triangular, but SO2 is angular; PCl5 is trigonal bipyramidal,
but SF4 is irregular tetrahedral; ClF3 is T-shaped, while XeF4 is
square planar.

Walsh diagrams propose a simple pictoral approach to determine
the geometry of a molecule considering and calculating the
energies of molecular orbitals of the molecule. The molecule will
have that geometry in which the energies of the molecular orbitals
used are minimum. Using this concept we can understand why H2O
is angular and BeH2 is linear; or why CH4 is tetrahedral and SF4 is
distorted tetrahedral.

Number of compounds of non-metallic elements of group VA and
VIA (N, P, O, S etc.) use d  -p  bond e.g. R3PO, H3SiOSiH3, N 
SF3, S4N4F4 etc.
The formation of d  -p  bond is common for all the second period
elements and is not important for the elements of third and higher
periods. The p  -d  bonding is more favorable than the d  -p 
bonding for higher atoms i.e. atoms of third and higher periods.

Bent's rule states that more electronegative substituents prefer
hybrid orbitals having less s-character and more electropositive
substituents prefer hybrid orbitals having more s-character. Thus in
40
CH2F2, the F-C-F bond angle is less than 109.5o indicating less
then 25% s-character in C-F bond, but the H-C-H bond angle is
larger indicating C-H bond has more than 25% s-character.

The energetics of hybridization also explains the geometry of a
molecule. The molecule will have the geometry which involves the
hybridization of lower energy. Thus while the lighter elements of
groups IIIA, IVA, VA and VIA use sp2, sp3, sp3d and sp3d2
hybridizations respectively for their hydride formation, the haviour
elements of these groups use their pure p-orbital, leaving a lp-sink
into a pure s-orbital. Hence in the heaviest elements of groups IIIA
and IVA (Tl, Sn and Pb) the stable oxidation states are two unit
less (1 and 2 respectively) then the maximum oxidation states (3
and 4 respectively) i.e. TlCl, SnCl2 and PbCl2. Similarly, the
hydrides of group VA and VIA elements are found to have bondangles considerably reduced as one moves down in these groups.

As the more electronegative atoms use those hybrid orbitals, which
have higher p-character and less electronegative atoms use hybrid
orbitals with more s-character the bond angle in many tetravalent
molecules is seen slightly distorted than 109o28', e.g. in CH3Cl,
HCH bond is 110o20'. This is known as Isovalent hybridization.

Apicophilicity
is
the
tendency
of
more
electronegative
substituents, in trigonal bipyramidal molecules (e.g. PCl4F or
PCl3F2) to seek out the apical orbitals, e.g. in oxysulfuranes.

The common reaction of covalently bonded molecules are atomic
inversion, Berrypseudorotation; nucleophilic displacement and free
radical mechanism. Atomic inversion is the simple inversion of
atoms about the central atom of the molecule. This is analogous to
41
the inversion of umbrella in a high wind. The energy barrier to
inversion is strongly dependent on the nature of the central atom
and that of substituents. Hence the separation of optical isomers
and their racemization is possible only in such case which have
values of energy barrier sufficiently high (e.g. methylpropylphenyl
phosphine, 120 KJ mol-1).

Berry pseudorotation involves scrambling of atoms at the
equatorial and axial position in a trigonal bipyramid geometry (e.g.
in PF5). This is believed to take place through conversion of the
ground state trigonal bipyramidal into a square pyramidal transition
state back to a new trigonal bipyramidal structure.

Nucleophilic displacement, e.g. the solvolysis (Alkoholysis) of a
chlorodialkyl phosphine oxide proceeds with inversion of
configuration of the phosphorus atom. The highly electronegative
entering and the leaving groups are favourably disposed at the axial
positions.
1.8
1.
CHECK YOUR PROGRESS: THE KEY
(a) (i)
(b)
2.
Linear
(ii)
Pyramidal
(iii)
Irregular tetrahedral
(iv)
Linear
+
NO2 > .NO2 > -NO2
(a) In H2O molecule, eight electrons are distributed such that
the configuration is 2  12 , 1  u2 , 1  2ux, 1  2uy, (i.e. 2a21, 2b22,
3a21, 1b21). Thus the energy of 1  ux (3a1) orbital is very
42
much less, indicating that the molecule is angular. Since in a
linear molecule 3a1 orbital remains non-bonding (i.e. has
very high energy).
(b) MgCl2 molecule.
3.
4.
(a) (i)
S4N4F4
(ii)
N  SF3
(iii)
H3SiOSiH3
(b) (i)
The dipole mement has low value (14.6 x 10-3 Cm).
(ii)
The bond dissociation energy is high (500-600 KJ).
(iii)
P-O bond length is short.
(a) (i)
Monovalency of Thallium (TlCl) and divalency of
Lead (PbCl2).
(ii)
Hybrids of VA and VI A group elements on moving
down the group.
(iii)
FCF bond angle is shorter than 109o28', but HCH
bond angle is larger than 109o28'.
(b) Apicophilicity is the tendency of more electro-negative
substitutes to seek out apical orbital in a triangular
pipyramidal (TBP) structure. It is well illustrated by
oxysulfuranes.
43
UNIT - 2 METAL LIGAND BONDING
Structure
2.0
Introduction
2.1
Objectives
2.2
Limitations of Crystal Field Theory (CFT)
2.3
Molecular Orbital Theory (MOT)
2.3.1 Octahedral complexes
2.3.2 Tetrahedral complexes
2.3.3 Square Planar complexes
2.3.4 similarity in CFT and MOT
2.4
Ligand Field Theory (LFT)
2.5
Let Us Sum Up
2.6
Check Your Progress : The Key
44
2.0
INTRODUCTION
The accidental discovery of hexaammine cobalt (III) chloride ,
CoCl36NH3 by Tassaert in 1799 raised the question how and why CoCl3
and NH3 each of which is stable compound , could combine to give yet
another very stable compound , CoCl3 - 6NH3 . The first satisfactory
explaination of the question was given by Werner in 1893 in terms of
primary and secondary , two types of valencies of metals . However it
was Sedgwick who in 1927 introducing the concept of coordinate - bond,
pointed out that Werner's primary-valency of metal ion is the
electrovalency, while the secondary valency is coordinate covalency , in
which the metal ion behaves as a 'Lewis - acid ' (electron pair acceptor).
Sidgwick suggested that the central atom or ion accepts n pairs (= its
coordination number) of electrons to attain effective atomic number
(EAN) , i e total number of electrons equal to the next noble gas. But the
exact nature of bonding and explaination of stereochemical, kinetic,
thermodynamic , spectral and magnetic properties could not be obtained
until the quantum mechanical theory was applied in the explaination of
complex formation.
Pauling's VBT enjoyed considerable support from 1930 to 1950 as
it explained the structural and magnetic properties of complexes in a
simple pictoral way involving hybridisation of orbitals on the metal ion.
Next came crystal field theory , an electrostatic approach, first
developed many years earlier in 1929 by Bethe and Van Vleck (1932).
Crystal field theory CFT) is concerned with the effect of the external
electric field due to the ligands on the relative energy levels of d-orbitals
of the central metal atom or ion in a regular octahedral field , for
example, the orbitals split into a group of three (t2g) of lower energy and
45
a group of two (eg) of higher energy. The nature of the splitting
determines the distribution of electrons among orbitals and hence leads to
the interpretation of magnetic and spectroscopic properties, and is often
very useful in the discussion of distorted structures, thermodynamic
spectral and magnetic properties. The MOT (Molecular Orbital Theory) is
most comprehensive of all the theories, was developed by Van Vleck
(1935) . CFT and MOT apparently give similar results. In practice, thus
modified CFT in which allowance is made for the overlap of the metal
and the ligand orbitals is preferred. The theory is then called 'the adjusted
crystal field theory' (ACFT) or the ligand field theory (LFT) in which the
simplicity of CFT is retained, but the M-L  and  -bonds are included .
Chemists have found that these two approaches are particularly
valuable for describing the bonding in transition-metal complexes : the
electrostatic "ligand-field" approach and the "molecular orbital" (MO)
approach. The methods of these theories are fundamentally quite different
; the first assumes essentially pure ionic interaction between the central
metal ion and the surrounding ligand atoms or anions, and the second
assumes the formation of covalent bonds between the central metal atom
and surrounding ligands. Yet as we shall see, these apparently
contradictory methods lead to electronic energy level diagrams which are
remarkably similar.
In this unit we describe both approaches, using only a few
examples to illustrate their application. You may recall what you have
already studied about the basic concepts of VBT, CFT and MOT as
applied to metal - ligand bond formation in coordination compounds .
46
2.1
OBJECTIVES
The main aim of this unit is to study nature of metal - ligand bond
and the mechanism of its formation. After going through this unit you
should be able to :
*
describe the limitations of crystal field theory (CFT);
*
discuss the mechanism of metal ligand bond - formation according
to MOT, for octahedral , tetrahedral and square planar complexes,
*
identify the similarities between CFT and MOT , and
*
discuss the ligand field Theory (LFT) of metal Ligand bond, as an
adjusted crystal field theory.
2.2
LIMITATIONS OF CRYSTAL FIELD THEORY
Crystal field theory (CFT) was developed (Bethe , 1929 and Van
Vleck , 1932) from 'electrostatic theory ' (Van Arkel , de Boer and Garric,
1932 ), which, considering ligands as point negative charge or point dipole works out the effect of the ligand field (Orderly arrangement of the
ligand around the central metal ion , similar to that in an ionic crystal) on
the electronic states of the metal. CFT does not consider mixing or
overlapping of ligand orbitals with those of metal ion , but calculates how
the repulsive effect of electronegative electrical potential field of the
ligands splits d - orbitals of the metal ion into groups of orbitals
(Crystal Field Splitting ) due to loss of the degeneracy .
Fig. 2.1 : Crystal field splitting of d - orbitals of central metal cation
of tetrahedral , octahedral , tetragonal and square planar complexes
47
According to this theory the physical and chemical properties of
complexes depend upon the energy of d - orbitals of metal ions , ie crystal
field splitting (CFS) and crystal field stabilisation energy (CFSE) . On the
basis of these only , this theory explains the bonding ability of ligands
(i.e. spectrochemical series) stability of metal complexes, their magnetic
properties (i.e. low-spin and high spin complexes); and the reactivity of
complexes and their thermodynamic properties . But this theory could not
explain the covalent character and the properties based on this character
of the coordinate - bond (i.e. a semipolar bond), thus:
(1)
The spectrochemical series:
There is a striking general pattern in the relative values of o for
different ligands. Almost irrespective of the nature of the metal ion
o , increases along the series (i.e. spectrochemical series):
However the theory presents no explanation for the relative
positions of H2O and -OH; F- and –CN and CO and C2O42-.
I-<Br-<Cl-<F-<OH-<C2O42-<H2O<Pyridine<NH3<ethylene
diammine< dipy <O-Phenan.<CN-<CO.
(2)
Nephelauxetic effect
The interelectronic repulsions as reflected by the actual spectra of
the complexes are less than those present in the free ions. This is
attributed to the effective increase in the size of the orbitals
housing the electrons by the combination of the orbitals of the
metal ion and those on the ligands giving larger room to the
elctrons to move around increasing , at the same time , the stability
of the complexes. This is termed as the nephelauxeticeffect (cloud
expansion effect) and is produced by the large sized ligands having
d or other suitable orbitals , like  * orbitals.
48
(3)
ESR , NQR and NMR spectra
ESR and NMR spectra of various complexes (IrCl62-, PtX42-, PdX42etc.) and NQR (Nuclear Quadrapol Resonance) spectra of CoCl42clearly indicate varying degree of covalency in these complexes ,
but CFT has no room for this character .
(4)
Radial Wave Function
Similarly the radial wave functions of the d orbitals and of the
ligands should have some overlap at the observed internuclear
distances in the metal complexes , ie the ligands are not point
charges , but have their own electron orbitals too. This too has no
explaination in CFT.
(5)
Bonding
The CFT can not account for bonding in complexes .
(6)
Charge Transfer Spectra
It does not explain the charge transfer (CT) spectra and the
intensities of the absorption bands.
(7)
Ferromagnetism and Antiferromagnetism
Neutron diffraction studies show the anti ferromagnetism to be due
to the tendency of half the ions to have their magnetic moments
aligned antiparallel to those of the other half of the ions due to the
overlap of the unpaired electrons on the metal with the orbital of
M2+—O2-—M2+. Because of the overlap of the opposite spin
electrons of the O2- ion with the two adjacent atoms , the two metal
atoms should also have opposite spins.
49
Check Your Progress - 1
Notes : (i) Write your answers in the space given below .
(ii)
(a)
Compare your answers with those given at the end of the
unit.
Proofs of covalency in complexes are presented by the evidences
such as
(i)
(ii)
(iii)
(b)
The anomaly in the positions of F and CN ions in the
spectrochemical series can be explained in terms of .
2.3
MOLECULAR ORBITAL THEORY (MOT)
The molecular orbital theory was developed by Hund and
Mulliken. According to Molecular orbital theory, which is modern and
more rational , a molecule is considered to be quite different from the
constituent atoms. All the electrons belonging to the atoms constituting a
molecule are considered to be moving along the entire molecule under the
influence of all the nuclei. Thus a molecule is supposed to have orbitals
of varying energy levels in the same way as an isolated atom has. These
orbitals are called molecular orbitals.
The molecular orbitals are filled in the same way as the atomic
orbitals , following the Aufbau and Pauli's Exclusion Principles. Thus , a
molecular orbital, like an atomic orbital , can contain a maximum number
of two electrons and the two electrons have opposite spins.
The difference between an atomic orbital and a molecular orbital is
that while an electron in a atomic orbital is influenced by one positive
50
nucleus, an electron in a molecular orbital is influenced by two or more
nuclei depending upon the number of atoms contained in a molecule.
Linear Combination of Atomic Orbitals (LCAO)
According to this approach , the molecular orbitals are formed by
the linear combination of atomic orbitals of the atoms which form that
molecule. For example, consider a hydrogen molecule constituting of two
atoms labelled as HA and HB. Let the wave functions in the atomic
orbitals of two atoms A and B be  A and  B respectively. When these
atomic orbitals are brought closer, they combine to form molecular
orbitals. In wave function  for electron in the field of nuclei of atoms A
and B may be written as suitable linear combination of wave function
 A and  B . As already discussed  may be expressed in two way .

=
A+ B
.........................................(2.1)
*
=
 A- B
.........................................(2.2)
Here  corresponds to wave function for symmetric combination
and is referred to as bonding molecular orbital, whereas  * corresponds
to anti symmetric combination and is referred to as antibonding molecular
orbital.
Now,  2 and ( *)2 give probabilities of finding electrons in the
two molecular orbitals formed according to equations (2.1) and (2.2). On
squaring equation (2.1) we get :

2
=
( A +  B)2
=
 A +  B + 2 A B
2
2
it is clear that  2 is greater than  A2 +  B2 by an amount
2 A B ie the probability of locating electrons in the molecular orbital
obtained by the linear combination in accordance with the equation (2.1)
is greater than that in either of the atomic orbitals  A and  B. In other
51
words, the molecular orbital represented by  has a lower energy than
either of the atomic orbitals represented by  A and  B . This orbital ,
therefore leads to the formation of a stable chemical bond and is,
therefore, termed as a 'bonding molecular orbital'.
On squaring equation (2.2) we get
( *)2 = ( A -  B)2 =  A2 +  B2 - 2 A B
Here ( *)2 is less than  A2 +  B2 . Hence the probability of
locating
electrons in the molecular orbital obtained by the linear
combination of atomic orbitals in accordance with equation (2.2) is less
than that in either of the atomic orbitals  A and  B. In other words, the
molecular orbital represented by  * in equation (2.2) has an energy
higher than the energy of either of the atomic orbitals represented by  A
and  B. Such an orbital, can not lead to the formation of a chemical
bond and is, therefore, called as an antibonding molecular orbital.
We may take a simple case of combination of 1s orbital of one
hydrogen atom with 1s orbital of another hydrogen atom to give two
MO's in an H2 molecule(Fig. 2.2).
Fig. 2.2: Formation of s bonding molecular orbital
The high electron concentration in between the two positive nuclei
shields them from mutual repulsion and holds them together at the
observed distance from each other. Such a molecular orbital is said to be
a bonding molecular orbital. It is designated as  (1s) orbital. The sign 
52
signifies that the orbital is symmetrical about the molecular axis and (1s)
indicates that it is formed by the combination of 1s atomic orbitals.
The second way of combing the two AOS is by subtraction, as
shown in the following figure (Fig. 2.3) :
Fig. 2.3 : Formation of s* antibonding molecular orbital
The two AOs being oriented in opposite directions cancel out so
that the probability of finding electrons in the region of overlap is
practically nil . The molecular orbital obtained in this manner is called an
antibonding molecular orbital.
A bonding MO has high electron density in the overlap region
whereas the antibonding MO has no electron density in between the
nuclei. This can be graphically illustrated by plotting the square of wave
function  2 (probability density) of orbitals of two H atoms. The
probability density for the 1s AOs of the individual atoms (i.e.  A2 and
2
 B ) are shown by dotted lines. When the probability densities of both
the atoms are added, the probability density function for the bonding MO
i.e.  2 MO is obtained. It is shown by solid lines. It is clear from the
following figure that there is greater probability of finding the electron
between the two nuclei than for separate atoms (Fig. 2.4).
Probability Density Plot
Fig. 2.4 : Sum combination of Wave-function
53
For antibonding MO, the picture is quite different. It is formed by
subtraction and he probability density of the two atomic orbitals get
cancelled in the centre so that there is no probability of finding the
electrons in the region overlap, i.e. between the two nuclei, as shown in
the following figure (2.5)
Zero Probability
Fig. 2.5 : Difference combination of wave-function
The lower energy of  (1s) orbinals and higher energy * (1s)
orbintal in hydrogen molecule is also represented graphically as follows.
It is the low energy of a bonding molecular orbital which makes it
bonding (Fig. 2.6).
Fig. 2.6 : Bonding and Antibonding MOs in H2
Similar to the formation of bicentric molecular orbital in H2
molecule , the formation of polycentric molecular orbitals in metalcomplexes
can be visualise ; e.g. the formation of bonding and
antibonding molecular orbitals as a result of overlapping of ns orbital of a
54
metal ion with the ligand group of orbitals matching symmetry in an
octahedral complex (Fig. 2.7).
Fig. – 2.7 : Formation of a polycentric molecular Orbital
The formation of molecular orbitals in a metal complex involves
the following steps :
1.
The first step is to identify the metal orbitals matching symmetry
for  bonding in a particular stereochemistry (i.e. geometry) of the
complex.
2.
The next step is to workout ligand group of orbitals (LGO) suitable
for overlapping with the metal atomic orbitals.
3.
Now, the bonding and the antibonding molecular orbitals are
formed as a result of the sum and the difference combinations of
wave functions of the metal and the LGOs (using LCAO method).
4.
In the next step, the  -molecular orbitals , thus formed in step -3,
are arranged in the order of increasing energy to give MO energylevel diagram.
5.
In the final step the total bonding electrons ( i.e. valency electrons
in the meta ion + the n number of electron - pairs obtained from n
ligands as a result of LM  -bond formation) are distributed in
the molecular orbitals according to aufbau and Hund rule.
2.3.1 Octahedral Complexes
In an octahedral complex, the metal ion is placed at the centre of
the octahedron and is surrounded by six ligands which reside at the six
corners of the octahedra (i.e. along +x , -x , +y, -y , +z and -z axes (Fig.
55
2.8). Using the steps given above the six metal - ligand  -molecular
orbitals can be constructed.
1.
Metal Orbitals suitable for a bonding.
The central metal caption of 3d-series elements contains in
all nine valence-shall orbitals which are: 4s, 4px, 4py, 4pz, 3dxy,
3dzx, 3dx2-v2 and 3dz2 .
All the nine atomic orbitals have been grouped into four
symmetry classes which are given below:
4s  A1g, or a1g; 4px, 4py, 4pz  T1u or t1u, 3dx2-y2
3dz2  Eg, or eg; and 3dxy, 3dyz, 3dzx  T2g or t2g
Now, we know that since in an octahedral complex six σorbitals of the six ligands are approaching along the axes (Fig. 2.8).
Fig. 2.8 Six ligand σ-orbitals in an octahedral complex.
in which ligands σ-orbitals (along the +x, -x, +y, -y, +z and -z axes
have been represented as σx, σ-x. σy, σ-y, σz and σ-z respectively).
These σ-orbitals, in order to form six metal-ligand σ-bonds or
MO's will overlap more effectively with only those metal ion
valence AO's that are having their lobes along the axes (i.e., along
the metal ligand direction). Quite evidently, such AO's are 4s, 4px,
4py, 4pz, 3dz-2 and 3dx2-y2, since these orbitals have their lobes lying
along the axes along which the six σ-orbitals of the six ligands are
56
approaching towards the central metal cation to form six metalligand σ-bonds.
The remaining three AO's namely 3dxy, 3dyz and 3dzx do not
participate in σ-bonding process, since these have their lobes
oriented in space between the axes. Thus these orbitals remain nonbonding and hence are called non-bonding orbitals.
2.
Construction of LGO for σ-bonding
The σ-orbitals of ligands combine together linearly to form
such group of ligand σ-orbitals that should be capable of
overlapping with the central metal ion six AO's viz 4s, 4p x, 4py,
4pz, 3dz2 and 3dx2-y2, e.g.(a) Since 4s orbitals has the same sign in all directions, the linear
combination of ligand σ-orbitals which can overlap with 4s orbitals
is: σx+ σ-x+ σy+ σ-y+ σz+d-z. This linear combination is represented
by  a which in its normalised form, is given by:
a 
1
 x  ó -x   y    y   z   z 
6
...............A1g or a1g
(A1g or a1g represents group symmetry class name of  a)
(b) Since one lobe of 4px orbitals has + sign and the other has - sign,
the linear combination, of ligand σ-orbitals that can overlap with
4pz orbital is σx-σ-x. It is represented by  x which in its normalised
form is given as:
a 
1
 x -   z 
2
..................Eg or eg
57
Similarly σy-σ-y and σz-σ-z are the linear combination of
ligand σ-orbitals that overlap with 4py and 4pz atomic orbitals
respectively. Thus:
y 
1
 y -   y 
2
..................Eg or eg
 z -   z 
..................Eg or eg
1
z 
2
(c) Since one opposite pair of lobes of 3dz2-y2 orbital has a + sign and
the other has - sign, the linear combination of ligand σ-orbitals for
this orbitals is: σx+ σ-x- σy σ-y. Thus

2
x y
2

1
 x    x    y    y 
2
..................E1u or e1u
(d) To find the ligand σ -orbitals combination for 3dz2 orbital poses
some difficulty. The analytical function for3dz2 orbitals is
proportional to 3z2-r2. The proper σ-orbital combination is easily
written down by substituting x2+y2+z2 for r2 in 3z2-r2. Thus 3z2r2=3z2-(x2+y2+z2)=2z2-x2-y2. Consequently the proper combination
for 3dz2 orbitals is:
2  z    z    x    x    y    y   2  z  2  z   x    x   y    y
Thus,
 z2 
1
2 3
2
z
 2ó -z  x  x   y    y 
..................T1u or t1u
(a) Overlap of metal 4s-orbital (a1g symmetry orbital) with a group
ligand σ-orbitals (a1g symmetry).
58
a 2 
1
σ x  ó- x  σ y  σ y  σz  σ z 
6
(b) Overlap of metal 4pz-orbital (t1u symmetry orbital) with a group
ligand σ-orbitals (t1u symmetry).
 x2 
1
 z -   z 
2
(c) Overlap of metal 3dz2-orbital (eg symmetry orbital) with a group
ligand σ-orbitals (eg symmetry).
 z2 
1
2 3
2
z
 2  z   x    x   y    y 
59
(d) Overlap of metal 3dz2-y2orbital (eg symmetry orbital) with a group
ligand σ-orbitals (eg symmetry).
 x 2  y2 
1
σ x  σ x  σ y  σ y 
2
Fig. 2.9 Construction of LGO's in octahedral complexes.
3. Formation of σ-molecular Orbital.
In the final step the six atomic orbitals of the central metal
cation viz 4s, 4px, 4py, 3dx2-y2 and 3dz2 overlap with six group
ligand σ-orbital viz  a,  x,  y,  z,  x2-y2 and  z2 respectively
to form six sigma bonding (abbreviated as σb) and six sigma
antibonding (abbreviated as σ*) molecular orbitals. Thus(a) 4s and  a which have the same symmetry (a1g symmetry) overlap
to form one σsb MO and one σs* MO.
(b) 4px and  x (both with eg symmetry) overlap to form one σxb -MO
and one σx* -MO.
(c) 4py and  y (both with eg symmetry) overlap to form one σyb -MO
and one σy* -MO.
(d) 4pz and  z (both with eg symmetry) overlap to generate one σzb MO and one σz* -MO.
(e) 3dz2 and  z2 (both with t1u symmetry) overlap to form one σz2.b
and one σz2* -MO.
(f) 3dx2-y2 and  x2-y2 (both with tu symmetry) overlap to form one σ
2 2bx -y MO
one σ x2-y2*-MO.
60
Thus we see that the combinations of six central metal
atomic orbitals with six ligand σ-orbitals give six σb and six σ*
molecular orbitals in an octahedral complex. It is thus obvious that
on adding twelve molecular orbitals (6σb- and 6σ*-MO's) to the
three non-bonding AO's viz. 3dxy, 3dyz, 3dzx we get in all fifteen
orbitals potentially available for electron filling.
4. Molecular Energy Diagram
The molecular orbitals thus formed in step 3, when arranged
in the order of increasing energy, we get molecular energy level
diagram. This can be obtained considering the following principles:
(i) Coulombic energies are in the order (legends are electronegative
than the metal ion): σligand < ligand <3d <4s <4p.
(ii) The mixing of the metal and the ligand group orbitals is
proportional to the overlap of the metal and ligand orbitals, and is
inversely proportional to their coulombic energy difference.
(iii) Bonding σ MOs are more stable than the bonding  orbital and
therefore, the antibonding σ*MOs are less stable than the
antibonding  * orbitals.
(iv) The bonding orbital lie closer to the LGOs, the antibonding
orbitals are closer to the metal orbitals due to the electronegativity
differences between the metal ion and the ligands atoms.
As the 4s and 4p orbitals can have a better overlap with
LGOs than the 3d orbitals can have with the LGOs, the a 1g and t1u,
MOs are at the lowest energy (corresponding a1g* and t1u* go to the
highest energy levels). The eg and eg* orbitals derived from the eg
orbitals of CFT, interact less with the LGOs due to a poorer
overlap, while the t2g orbitals remain nonbonding in a σ only
picture and are not displaced. The energy level diagram is given in
Fig. 2.10 for the octahedral complexes.
61
Fig. 2.10: MO energy level diagram for an octahedral
complex with no  bonding.
As can be seen in Fig. 2.10, the 15 molecular orbitals are
arranged in 7 energy levels. As the energies of all the orbitals in a
symmetry group are equal (i.e. degenerated), out of these seven
energy levels, three are triply degenerated (i.e. T1u, T2g and T1*u).
two are doubly degenerated (i.e. Eg and Eg*) and the remaining
two are monodegenerated (i.e. Alg and A*lg). These energy levels
are arranged as follows:
A1g < T1u < Eg
Bonding MO's
<
T2g
Non-bonding MO's
< Eg* < A*lg < T1u*
Antibonding MO's
As the ligand orbitals (shown on the right hand side in the
Fig. 2.10) are more negative than the metal orbitals (shown on the
left hand side in the Fig.), hence they are shown relatively below
(of lower energy) than the metal orbitals. Thus their energy is
comparable with the bonding molecular orbitals.
62
5.
Distribution of electrons in the Molecular Orbitals
As has been mentioned above if a MO is near in energy to
the energy of the ligand orbitals, it would have more of the
character of the ligand. Thus six σb-MO's (i.e. σsb, σx2, σy2, σzb σx2-y2
and σz2 MO's) which are nearer the energy level of the ligands, are
occupied mainly by the ligand electrons. In other words the
electrons in the six σb-MO's are mainly localised on ligand orbitals,
since σ-orbitals of the ligands are more stable than the metal
orbitals. Conversely, electrons occupying any of the six σ *-MO's
are to be considered mainly metal ion electrons. Electrons in t 2g set
of orbitals will be purery metal ion electrons when there are no
ligand  orbital.
Further, the crystal field splitting energy (  o or 10 Dq) in an
octahedral complex, according to MOT, is the difference in energy
between the t2g(dxy,dyz, and dzx orbitals) and eg*(σx2-y2* and σz2*
MO's) energy levels.
In case of weaker ligands such as F-ion, the energy
difference,  o between the t2g set eg* set is smaller than P (i.e.
 o<p) and hence the lowest-energy antibonding MO's namely σx
2
y *
2
-
and σz2* have approximately the same energy as the non-
bonding AO's: 3dxy, 3dyz and 3dzx (t2g set). Consequently spin free
(or high spin) complexes are formed. However, the strong(er)
ligands such as NH3 molecules split the σ-bonding MO's more
widely and the energy difference,  o between the t2g-set of nonbonding AO's (3dxy, 3dyz and 3dzx) and eg*-set of MO's (σx2-y2* and
σg2*MO's) is greater than the electron pairing energy. P (i.e.  o >
P). Hence spin paired (or low spin) complexes are formed. For
example Co(III) complexes in an decahedral stereochemistry, in
63
all have 18 valence electrons, 12 electrons from six ligands and 6
from d-orbitals in Co3+ ion.
The distribution of 18 electrons in [CoF6]3- which contains
weak (er) ligands viz. F ions (i.e. it is a high-spin complex) takes
place in various orbitals according to the above scheme will be as
follows:
(σs)2, (σx)2 = (σy)2 = (σz)2, σ(x2-y2)2 = (σz2)2, (3dx)2 = (3dyz)1 =
(3dxz)1 = σ*(x2-y2)1 = σ*(z2)1
Hence it will be paramagnetic due to four unpaired electrons.
While, the distribution of electrons, in [Co(NH3)6]3+ which contains
strong(er) ligans viz. NH3 molecules (i.e. it is a low-spin complex)
takes place in various orbitals as follows:
(σs)2, (σx)2 = (σy)2 = (σz)2, σ(x2-y2)2 = (σz2)2, (3dxy)2 = (3dyz)2 =
(3dxz)2
Thus the complex is diamagnetic, due to all paired electrons.
2.3.2 Tetrahedral Complexes
The method used for the construction of molecular orbital energy
level diagram for the octahedral complex, may also be used for the
construction of molecular energy diagram for tetrahedral complexes. Out
of the nine valency orbitals of the metal ion, the orbitals matching to the
Td-Symmetry are s (a1g) and p (t1u). In addition to this, out of the five dorbitals, three, the t2g (dxy, dxz and dyz) orbitals are also suitable for the σbonding. Thus the combined symmetry of p(t1u) and d(t2g) orbitals is
generally expressed by T2(Cf the character table for Td symmetry).
Similarly, on working for the ligand groups of orbitals (LGO), the
ligand orbitals with a pair of electrons in each, give one LGO of T2
symmetry,  t2, and one LGO of a1 symmetry. The LGO of T2 symmetry
will overlap metal ion orbitals of both the groups (p and d, i.e. t1u and t2g).
64
As a result, three triply degenerated energy levels of molecular orbitals
will be formed; one the group of bonding molecular orbitals, the other
one, the group of slightly anti-bonding molecular orbital and the third one
the group of completely anti-bonding molecular orbitals. (Fig. 2.11).
It may be mentioned, contrary to the octahedral complexes, here
(in the tetrahedral complexes) eg orbitals are the non-bonding orbitals;
and the energy difference between the eg and the next higher t2* energy,
levels is  t (Analogous to CFT).
Fig. 2.11: Molecular energy level diagram for a tetrahedral complex.
Now the electrons are distributed in these molecular orbitals
following the aufbau principal and Hand rule. Thus in [CoCl4]2-, there are
15 valency electrons (8 from the four ligands + 7 electrons of Co(II) ion).
Out of these, 12 electrons are used to saturate six bonding molecular
orbitals (i.e. t2, a1g and eg orbitals), while the remaining three electrons are
present one each in the slightly anti-bonding, three t*2 group orbitals (As
Cl- is a weak ligand hence  t <P). Hence [CoCl4]2- complex is
paramagnetic due to three unpaired electrons.
65
2.3.3 Square Planar Complexes
The square planar complexes, ML4 have D4h symmetry. In these
complexes five d-orbitals of the metal ion, loosing their degeneracy, split
into the four groups of symmetry a mono degenerate a 1g (dz2), a doubly
degenerate eg (dxy and dyz) and the remaining two mono-degenerate
b2g(dxy) and b1g(dx2-y2) orbitals. Similarly, metal ion p(t1u) orbitals also
loose their degeneracy and split into two groups, a monodegenerate
a2u(pz) and a doubly degenerate (px and py) symmetry orbitals.
The four ligands present on the two x and y axes give ligand group
of orbitals (LGO) of a2g, b1g and eu symmetry orbitals. (Fig. 2.12)
Fig. 2.12: LGO's in a square planar complex.
These LGO's overlap with the metal ion-orbitals of same symmetry
and give same number of bonding and anti-bonding molecular orbitals.
While, the a2u, eg and b2g symmetry orbitals remain non-bonding in the
complex. (Fig. 2.13)
66
Fig. 2.13: Molecular Orbital Energy Level Diagram in a
Square Planar Complex.
67
Check Your Progress - 2
Notes :
(i)
Write your answers in the space given below .
(ii)
Compare your answers with those given at the end of
the unit.
(a)
Name the metal orbitals and their symmetry groups suitable for σbonding in metal complexes of following stereochemistry along
with the LGO's of matching symmetry:
Metal σ-orbitals
Symmetry
LGO
(i)
Octahedral ............................. ............................. ..........................
(ii)
Tetrahedral ............................. ............................. ..........................
(iii) Square
planar
(b)
............................. ............................. ..........................
 E (cf crystal field splitting in CFT) is the difference between the
following
molecular
orbital
energy
levels
in
the
given
stereochemistry of metal complexes:
(i)
Octahedral
...........................................
(ii)
Tetrahedral
...........................................
(iii)
Square planar
...........................................
2.3.4 Similarity in CFT and MOT
Thus, we have seen both CFT and MOT consider that the
degeneracy of metal d-orbtials is lost after complex formation; e.g.
according to CFT in the octahedral complexes metal d-orbitals split into
tzg and eg groups. The energy of eg group being higher than that of t2g
group (Fig. 2.1). Similarly, according to MOT, after σ-overlap of metal
orbital with ligand group orbitals (LGO) in the octahedral stereochemistry
68
the t2g symmetry orbitals form the non-bonding energy level in the
complex, while the anti-bonding eg, E*g, form the higher energy level
above the T2g level. (Fig. 2.10)
Further, similar to CFT, in MOT also, the energy difference
between t2g and eg levels (E*g level in MOT) is taken as  o (crystal field
splitting, CFS, in CFT). Which forms the basis for the explanation of
most of the properties of complexes.
Thus the results of both CFT and MOT are same as far as dbonding is concerned. Only the mechanisms are different in these two
theories. While the cause of  o (CFS) in CFT is the electrostatic
repulsion, according to MOT this results overlap of the metal orbitals
with the ligand orbitals. There has, of course, been no change in  o which
is any way an experimental quantity; all that has changed is the
interpretation of it.
2.4
LIGNAD FILED THEORY (LFT)
2.4.1 -Bonding and Molecular Orbital Theory
As has been mentioned earlier, if the ligands have -orbitals of the
correct symmetry, filled or unfilled, their interaction with the -orbitals of
metal is considered. For the  -bonding the ligand should be a very good
Lewis-base, but generally it does not happen so. There are many common
ligands, which are very weak electron-doners, still they form stable
complexes. e.g. CO, RNC, PX3 (x=Halogen), PR3, AsR3, SR2, C2H4 etc.
As a matter fact  -bonding plays an important part in their stability.  bonding may be of four types:
69
1.
M(d  )  L(p  ) bonding
In which d  electrons of the metal ion are donated to the
vacant p  orbitals of the ligand (i.e. back bonding). Such bonds
are formed by the metals, situated towards the end of a transition
series in their lower oxidation states with the unsaturated ligands
e.g. CO, NO2-, CN-, RNC, NO etc. (Fig. 2.14)
2.
M(d  )  L(d  ) bonding
In which d  electrons of the metal ion are donated to the
vacant d  orbitals of the ligand. Such  -bonds are formed by the
metals, placed towards the end of a transition series in their lower
oxidation states with the ligands e.g. P, As, S etc. as the acceptor
atoms (e.g. PF3, PR3, AsR3, SR2 etc.) (Fig. 2.15).
3.
L(p  )  M(p  ) bonding
Such  -bonding is rare, and is seen only in Be and B with
the ligands. Such as O2-, F-, NH2 etc. In this type of bonding p 
electrons of the ligand are given to the vacant p  orbitals of the
metal ion. (Fig. 2.16).
4.
L(p  )  M(d  ) bonding
This type of  -bonding is seen in the metals, placed in the
beginning of a transition series, in their higher oxidation states with
the saturated ligands such as O2-, F-, NH2- etc. (Fig. 2.17).
Fig. 2.14: M(d  )  L(p  ) bonding
70
Fig. 2.15: M(d  )  L(d  ) bonding
Fig. 2.16: L(p  )  M(p  ) bonding
Fig. 2.17: L(p  )  M(d  ) bonding
2.4.2 Formation of  - Molecular Orbitals:  -Bonding in Octahedral
Complexes:
Construction of  -molecular orbitals follows the first three
steps, used in the formation of σ-molecular orbitals; e.g. (i) Selection of
metal  - orbitals
(ii) Construction LGO's of matching symmetry and
(iii) formation of  -bonding molecular orbitals as a result of the overlap
of  -metal orbitals with the  LGO's.
71
Remaining two steps involve in the adjustment of  -bonding and
anti-bonding molecular orbitals in the σ-molecular orbitals energy level
diagram (Fig. 2.10) and (v) the redistribution of electrons in the modified
MO energy level diagram after adjusting  -bonding and anti-bonding
molecular orbitals.
1.
Metal  Orbitals
In the Octahedral complexes, the metal orbitals suitable for
 -bonding may be of t2g or t1u symmetry. As the orbitals of t1u
symmetry (i.e. p orbitals) if used for  -bonding will weaken  bonding, hence the t2g orbitals of metals are most suitable for M 
L  -bonding.
2.
Ligand  Orbitals
For formation of the  -bond, the ligand can use (i) a porbital perpendicular to the σ-bond axis (e.g. Cl-, F-, OH-) or
(ii) a d orbitals (e.g. PH3, PR3, AsR3 etc.), or
(iii) an anti-bonding  * molecular orbital (e.g. CO, CN-, Py)
3.
Formation of  LGO's
When the ligand  -orbitals is p  (p) or d  , based on the
symmetry of the metal ion  orbital (i.e. t2g), ligand group orbitals
may easily be constructed e.g. in Md  -Lp  bonding:
We know in an octahedral complex each ligand has two porbitals suitable for  -bonding (Fig. 2.18) e.g.on x axis
In ligand L1
 1x,  1y
on -x axis
In ligand L2
 2x,  2y
on y axis
In ligand L3
 3z,  3x
on -y axis
In ligand L4
 4z,  4x
on z axis
In ligand L5
 5x,  5y
on -z axis
In ligand L6
 6x,  6y
72
In Fig. 2.19 fot Mp  -Lp  bonding the overlap of pz orbital
of the metal cation with ligand group p  orbitals viz ½ (  y1+  x2 x3-  x4) is shown.
Fig. 2.18: p  orbitals of ligands in an octahedral stereochemistry
Fig. 2.19: Ligand group of orbitals matching symmetry with pz
orbital of the metal ion.
Similarly, for Md  - Lp  bonding, overlaps of metal dxz
orbitals with the ligand group of p  orbitals is shown (Fig. 2.20).
In table 2.1 are given ligand group ortitals (LGO) matching
symmetry with metal orbitals of t1u and t2g symmetries.
73
Fig. 2.20: LGOs for dxz metal  orbitals.
Table 2.1: LGO's matching symmetry with metal t1u and t2g
group orbitals for  bonding is an octahedral complex:
Symbol Metal Orbital
t1u
t2g
4.
Combination of the ligand
orbitals (LGOs)












px
1
 y 2   x5   x 4   y 6
2
py
1
 x1   y 5   y 3   x 6
2
pz
1
 y1   x 2   x3   y 4
2
pzx
1
 x5   y1   x 3   y 6
2
dyz
1
 x2   y5   x6   y 4
2
dxy
1
 x1   y 2   x3   y 4
2
Formation of  bonding and anti-bonding MOs
Similar to the  -bonding,  -bonding and anti-bonding
molecular orbitals are formed as a result of positive and negative
overlap of the wave functions of  -metal orbitals (t2g orbitals) with
the LOGs of same symmetry (  t 2 g ):
74
Bonding  -molecular orbitals:  t2g = T t 2 g   t 2 g 
Antibonding  -molecular orbitals:  *t2g = T t 2 g   t 2 g 
It is important to note that both the bonding and anti-bonding
molecular orbitals will be triply degenerated.
5.
Accommodation of  molecular Orbitals in the  molecular
orbital energy level diagram
The energy of  -molecular orbitals depends on two factors:
(i)
as compared to the energy of metal ion  -orbitals (t2g),
whether the energy of LOGs t2g symmetry is lower or higher;
and
(ii)
whether the ligand  -orbitals are saturated or vacant ?
This information is important in the light of the fact that
during combination of the orbitals the energy of orbitals having
lower energy is further lowered, while that of higher energy is
raised further. Hence;
(a)
when the ligand orbitals are vacant and are of higher energy
than that of the filled metal orbitals, M  L,  -bonding
takes place; but
(b)
when the ligand orbitals are saturated and are of lower
energy than that of the metal orbitals, L  M,  -bonding
takes place.
Metal ions placed towards the end of a transition series in
their lower oxidation state give the first type of  -bonding (i.e. M
 L,  -bonding) with the unsaturated ligands, such as phophine,
arsine, CN-, CO, NO etc. This type of  -bond increases stability of
the complex.
75
The other type of  -bonds are formed by the metal ions
placed at the beginning of a transition series in their higher
oxidation states with the saturated ligands such as F-, Cl-, O--, OHetc. This type of  -bonding destabilise metal complexes.
As a result or  -bonding, in  -molecular energy level
diagram (2.10), in place of non-bonding t2g-orbitals, t2g  -bonding
and t2g  * anti-bonding molecular orbitals are accommodated.
When there is M  L,  -bonding, the bonding  t2g molecular
orbitals (triply degenerated) find their place above the e g group of
orbitals, while the anti-bonding  *t2g orbitals get their position
above the anti-bonding eg*, but below the anti-bonding a1g*
orbitals (Fig. 2.21)
Similarly, in case of L  M,  -bonding, the triply
degenerate bonding  t2g orbitals are placed above the bonding egmolecular orbitals, while the anti-bonding  - t2g* orbitals find their
place below the anti-bonding eg*-molecular orbitals (Fig. 2.22):
Fig. 2.21: M  L,  -bonding.
76
Fig. 2.22: L  M,  -bonding.
2.4.2 Effect of  -bonding on ligand field splitting energy  o
As due to  -bonding the energy of non-bonding t2g molecular
orbitals is changed, the value of ligand field splitting energy is also
changed. While the M  L,  -bonding increases the value of  o, the L 
M,  -bonding very much decreases it. (Fig 2.21 and 2.22 respectively).
M  L,  -bonding:
In this case the net result of  -interaction is that the metal t2gorbitals are stabilised relative to the eg*-MOs i.e. the metal t2g electrons
will go into the t2gb-MOs which are of lower energy than t2g*-MO's and
the thus the value of  o will be increased to  o (  1o<  1o). In this case
ligand exerts a stronger field. A ligand of this type is referred to as an
acceptor ligand because of the presence of empty  -orbitals in it and the
 -bonding established in such a case is sometimes referred to as metal-
to-ligand (M  L)  -bonding. Phosphines arsines and CO are important
examples of this type of ligands. Hence their complexes are very stable.
L M,  -bonding:
In this case the t2g metal orbitals are de-stabilised relative to egMOS. The electrons of the ligand  -orbitals enter the lower t2g*-MOs and
those of the t2g metal orbitals will go to higher t2g* MOs. Since in this
77
case, the  -interaction destabilises the t2g-metal orbitals relative to eg*MOs, the value of  o is diminished to  o as shown in Fig. 2.22 (  o>  o).
The ligand in this case exerts a weaker field. A ligand of this type is
generally called a donor ligand because of its filled  -orbitals. Halide
ions are important examples of this type of ligands and the  -bonding of
this type is generally referred to as ligand-to-metal (LM)  -bonding.
Such type of  -bonding occurs in complexes having metal ions in their
normal oxidation states (especially lower oxidation states). These
complexes are comparatively less stable.
2.4.3 Ligand field theory: An adjusted crystal field theory (ACFT)
As has been pointed out earlier, ligand field theory was specially
developed to explain the nature of metal ligand bond in metal complexes.
We know, the crystal field theory takes no account of possible covalent
bonding in complexes and regards the bonding as purely electrostatic. But
the physical measurements such as electron spin resonance. NMR and
nuclear quadruple resonance suggest that there is some measure of
covalent bonding also in complexes. It is because of this reason that a
kind of modified form of CFT has been suggested in which some
parameters are empirically adjusted to allow for covalence in complexes
without explicitly introducing covalence into CFT. This modified form of
CFT is often called 'Ligand Field Theory' (LFT).
Ligand field theory, which is a combination of crystal field theory
and molecular orbital theory, is therefore in principle the most
satisfactory of the theories of bonding discussed in this unit. It may be
considered as the combination of  -bonding in CFT or we may say LFT
calculates the effect of  -bonding on the crystal field splitting. Thus,
LFT explains the mechanism of metal-ligand bonding in two steps . The
78
first step is to workout splitting of metal d-orbitals into groups of orbitals
having different energies in symmetry. In the second step it calculates the
effect of  -bonding (if it takes place) on the crystal field splitting and the
crystal field stabilization energy (CFSE). Probably, because of this
adjustment, Cotton and Wilkinson have called LFT, 'Adjusted Crystal
Field Theory' (ACFT).
Advantages of LFT
1.
As LFT was developed specially to explain metal ligand coordinate
bonding in complexes. Hence it is most successful theory.
2.
It has room for both electrovalence and covalence hence it explains
the properties of metal complexes more efficiently.
3.
It explains the positions of ligands in the spectrochemical series
satisfactorily. It should be noted that since it is an experimental
series it incorporates all effects of the lignads in splitting the dorbitals (including  -bonding). Hence, the anomalies of CFT,
failing in the explanation of the relative positions of F- and CN-,
and OH- and H2O in the spectrochemical series, can be explained
by LFT satisfactorily. Since CN- form ML  -bond giving large
values of  o, hence placed at the top of the series; but the saturated
ligands like F- form LM  -bond resulting in very small value of
 o, justifying its position at the base of the series.
4.
LFT also explains charge transfer bonds (cf MOT).
79
Check Your Progress - 3
Notes :
(i)
Write your answers in the space given below .
(ii)
Compare your answers with those given at the end of
the unit.
(a)(i) Which type of  -bonding increases the stability of a metal
complexes? (ii) and why? (iii) Give one example?
(b)
(i)
...........................................
(ii)
...........................................
(iii)
...........................................
Give (i) the names and (ii) the symmetry groups of metal  orbitals; and also (iii) the combination of ligand orbitals (LGO)
matching the symmetry in an octahedral complex.
(i) Names
(ii) Symmetry
(iii) LGO
.............................
.............................
..........................
.............................
.............................
..........................
.............................
.............................
..........................
2.5
LET US SUM UP

Although Werner (1893) proposed two types of valencies (Primary
and Secondary) for metal ions in, so called the compounds of
higher order (i.e. metal complexes) and Sidgwick (1927)
introducing the concept of coordinate bond and EAN explained
primary valency as electrovalency and the secondary valency as
coordinate covalency. The exact nature of M-L bond in the
complex compounds could not be understood, until the quantum
mechanical theory was used.
80

Out of the four theories proposed, based on quantum mechanics,
the two, VBT and MOT were covalent models, while CFT
considered M-L bond purely electrostatic, similar to that existing in
ionic crystals. However, experimental proofs (NMR, ESR and
Nephclauxetic studies etc) clearly indicated varying degree of
covalence in the M-L bond. Hence, LFT was formulated using the
two most useful theories CFT and MOT. Since the results of CFT
and  -bonding in MOT are same in determining splitting of metal
d-orbitals; and as the crystal field splitting energy,  E, form the
basis for the explanation of most of the properties of complexes,
LFT was developed by adding  -bonding from MOT in CFT, to
account for the covalence in M-L bond. Cotton and Wilkinson
called it 'Adjusted Crystal Field Theory' (ACFT).

The steps involved in the construction of molecular orbital in a
metal complex, are- (i) identification of metal and ligand orbitals
suitable for σ and  -bonding; (ii) formation of LGO to make
effective overlap with the metal orbitals using LCAO method.
(iii) Construction of bonding and anti-bonding molecular orbitals
as a result of sum and difference combination of the metal orbitals
with the LGOs of same symmetry for σ-bonding; and arranging
them to get molecular orbital energy level diagram; (iv) Now,  molecular orbitals (bonding and anti-bonding) are constructed, and
are adjusted in the energy level diagram prepared for σ-bonding;
and (v) the last step is to distribute total valency electrons (metal
electrons+those obtained from ligands as a result of donation) in
the molecular energy level diagram. The bonding
molecular
orbitals get the ligand-electrons; while the metal-electrons are
accommodated in the non-bonding and anti-bonding MOs.
81
Metal complexes may have either ML  -bonding or LM  -

bonding. The first one stabilities the complex, by increasing the
value of  E; while the later one destabilize it, as the value of  E
decreases.
In Oh complexes metal  -orbitals are t2g orbitals, while the lignad

may use p  , d  or anti-bonding molecular orbital for  -bonding.

LET has been found most satisfactory for metal complexes, as it
has room for both electrovalence and covalence, suitable for the
semipolar nature of coordinate bond. Further, it can explain the
positions of CN- and F- or HO- and H2O in the spectrochemical
series satisfactorily, which were otherwise considered anomalous
according to CFT.
2.6
CHECK YOUR PROGRESS: THE KEY
1.(a) (i) ESR (ii) NMR spectra and (iii) Nephelauxetic studies.
(b)  -bonding; since ML  -bonding in CN- complexes increases
the value of  E, Hence it is at the top of spectrochemical series
(SCS); while LM  -bonding in F- complexes very much
reduces  E, placing F- at the bottom in SCS.
2. (a)
(i) Octahedral
Metal σ -orbitals
Symmetry
LGO
s and p; dz2 and dx2-y2
a1g and t1u ,eg
Ʃa, Ʃt1u
and Ʃeg
Ʃa, Ʃt2g
(ii) Tetrahedral
s, dxy, dyz and dxz
a1g, t2g
(iii)
dx2-y2,
b1g,  1g , b2g, Ʃb1, Ʃay
planar
Square
2
dz , dxy, dyz
and dxz
82
eg
Ʃb2 Ʃeg
(b) (i)
Eg* - t2g
(ii)
t2* - Eg
(iii)
B1g* - A1g
3.(a) (i)
ML  -bonding
(ii)
It increase Crystal Field Splitting Energy  E very much.
(iii)
[Fe(CN)6]4-
(b)
(i) Names
(ii) Symmetry
(iii) LGO

dxy
1
 x1   y 2   x3   y 4
2
dxz
t2g

1
 x5   y1   x3   x6 
2
1
 x 2   y 5   x 6   y 4 
2
dyz
83
UNIT 3
ELECTRONIC SPECTRA OF TRANSITION METAL
COMPLEXES
Structure
3.0
Introduction
3.1
Objectives
3.2
Types of Spectra
3.3
Spectroscopic Ground States
3.3.1 Correlation
3.3.2 Selection Rules
3.4
Orgel Diagrams
3.5
Tanabe Sugano Diagrams
3.6.
Calculations of Dq, B and  Parameters
3.7
Charge Transfer Spectra
3.8
Spectroscopic method of assignments of absolute configuration in
optically
active
metal
chelates
and
information.
3.8.1 Optical Rotatory Dispersion, ORD.
3.8.2 Circular Dichroism, CD.
3.9
Let us sum up
3.10 Check Your Progress : The Key
84
their
stereochemical
3.0
INTRODUCTION
Explanation of electronic spectra of transition metal complexes,
could become possible only after developments of CFT. Which
considered splitting of a d-atomic orbitals of transition metal ions, after
complex - formation. This splitting, and hence splitting energy was
considered responsible for the visible and ultra-violet spectra of transition
metal complexes. However, the detailed explanation of transition metal
complex-spectra is obtained, with the knowledge of molecular orbitals of
metal complexes. The bands obtained in the electronic spectra of
transition metal complexes are considered to be due to transition of
electrons from one d atomic orbital to the other d
orbital (d – d
transitions). Many of the topics treated here are elaborations of the
material introduced in unit 2. However, we should draw attention to the
fact not all bands in the visible and U.V. spectra of transition metal
complexes are d – d (or Ligand field) spectra. Some arise from electron
transfer (which may be in either direction) between metal ion and ligand,
and are called charge transfer spectra.
The principal new work in the first part of this unit is the problem
of interelctronic repulsions. To see how to take repulsion into account,
you must recall what you have already studied about the basic concept of
interelectronic repulsion for atomic spectra, quantitatively. Then we can
go on to see how repulsions compete with the ligand field when the atoms
and ions are present as a part of a complex.
In the end of the unit we shall try to use these spectra for
assignments of absolute configuration in optically active metal chelates
and their stereochemical information.
3.1
OBJECTIVES :
85
The main aim of this unit is to see how to analyse the electronic
spectra of transition metal complexes, and hence to enrich our
understanding of their bonding. After going through this unit you should
be able to :

describe spectroscopic ground states and their correlation;

explain selection rules for d-d transition;

discuss orgel and Tanabe-Sugano diagrams for splitting of
electronic states in different ligand – fields;

calculate values of dq, B and  parameter considering the bands
obtained in the spectra of the complex;

explain charge transfer spectra (both the metal to ligand, and ligand
to metal charge transfer); and

assign absolute configuration of optically active metal chelates
using
'Optical
Rotatory
Dispersion'
(ORD)
and
'Circular
Dichroism' (CD).
3.2
TYPES OF SPECTRA
Most of the transition metal complexes are highly coloured and
absorb radiant energy in the visible region of the spectrum. As most of
the transitions in the visible spectrum involve the electron transitions
from one level to the other, it becomes necessary to study the population
of the energy levels at the ground state and the excited states, and the
probability of occurrence of the different transitions.
It should be mentioned that not all bands in the visible and
ultraviolet spectra of transition metal complexes are d – d spectra. Some
arise from electron transfer (may be in either direction) between metal ion
and ligand. Such charge transfer spectra are of high intensity and usually
86
occur at some what higher frequency than d-d transition. Generally,
electronic spectra are of four types :
(1)
The d-d or Ligand Field Spectra
This occurs in the near infrared, visible, and ultraviolet regions
(10000-30000 cm-1, 1000-333 nm). Lower frequencies are not
accessible
experimentally;
the
higher
frequencies
though
accessible, are overshadowed by the charge transfer and the
interligand transitions. This limits the study of the d-d transitions to
only the visible regions of spectrum. These transitions are
considered to be totally within the metal ion in the CFT model,
though some ligand contribution is included in LFT or ACFT
models. MOT treats these transitions as arising due to the
excitation of the electron from the t2g level to eg* levels belonging
largely to the metal itself.
(2)
Ligand-to-Metal Charge Transfer Bands
When the electron transition takes place from a MO located
primarily on the ligand (M – L bonding  or  orbitals) to a
nonbonding or antibonding MO located primarily on the metal
atom, the lgand-to-metal charge transfer bands are observed. These
cannot be explained by CFT and represent the tendency of ligands
to reduce the metal ion. The semiempiral MOT is adequate for
explaining it.
(3)
Metal-to-Ligand Charge Transfer Bands
These involve the transition of electron from an antibonding or
non-bonding orbital, concentrated on the metal atom, to the
antibonding orbital located primarily on the ligand, and measures
the tendency of the metal ion to reduce the ligand. These bands are
87
observed generally for the metal ions in low oxidation states in the
ultraviolet region, but are seen many times to tail into the visible
regions, e.g. [Fe (dipy)3]2+.
(4)
The Intraligand Transitions
When an electron transition takes place from one ligand orbital to
another ligand orbital, the intraligand transitions are observed.
They are found in the ultraviolet regions and can be readily
separated from the equally intense M – L charge transfer bands as
they are not affected much by the other ligands. They, however,
depend on the M – L bond strength.
3.3
SPECTROSCOPIC GROUND STATES
In absence of an external field, the five d orbitals are degenerate
and the ground state of a dn is given in Table 3.2. In the absence of the
external field, in fact the five d orbitals remain undefined. Under the
influence of an octahedral field, these split into a group of t 2g and eg (or
eg*) orbitals. An s orbital, being spherically symmetrical and nondegenerate, is not split into any states at all, and the p orbitals are not split
in octahedral field as they interact with the field to the same extent. f
orbitals get split into three levels – a triply degenerate t1 level, a triply
degenerate t2 level and a singlet level a2 (Fig. 3.1).
88
3.3.1 Correlation Diagrams
To develop energy level diagrams of complexes, start is made with
the free ion of dn configuration and then interelectronic repulsions and the
ligand field effects are added to it. Electron – electron repulsion is the
cause of splitting in the terms of electron configuration.
In the absence of an external field, a d1 configuration becomes 2D
which breaks into two energy levels, 2T2 (corresponding to t2g) and 2Eg
(corresponding to the eg) levels. A d2 configuration has two Russell –
Saunders states : a low energy 3F and the high energy 3P state. These
states behave, in an octahedral field, exactly in the same way as the f and
p orbitals.
Hence, a 3F state gives 3T1g, 3T2g and 3A2g states whereas the 3P
state is not split and gives 3T1g (P) state (Fig. 3.2). In dn configuration
electron – electron interaction starts.
89
Table 3.1 Splitting of the Terms in a dn ion in Weak Octahedral Field
Configuration
of free ion
d1
Ground
state of
free ion
2
D
Energy
level
diagram
A
Configuration Ground
of free ion
state of
the ion
6
5
d
D
Energy
level
diagram
A
d2
3
F
B
d7
4
F
B
d3
4
F
inverted B
d8
3
F
inverted B
d4
5
D
inverted A
d9
2
d5
6
no splitting
d10
1
S
D
inverted A
no spliting
S
* With reference to Fig. 3.2
Note that the ground states of a dn configuration involve S, P, D
and F states only (Table 3.1). This S and P states are not split in an
octahedral field, D and F states split as shown in Fig. 3.2. By considering
the hole formalism, all the splitting, can be explained for a dn system in
terms of Fig. 3.2 as A, B, inverted A or inverted B (Table 3.2, Fig. 3.3).
Table 3.2 Splitting of the Ground State Terms in an Octahedral Field
Term
S
Components in octahedral field
A1g
P
T1g
D
Eg + T2g
F
A2g + T2g + T1g
G
A1g + Eg + T1g + T2g
H
Eg + Tg + T1g + T2g
I
A1g + A2g + Eg + T1g + T2g + T2g
90
Fig. 3.3 Ground term splitting and energy (in Dq) unit for dn ions in
octahedral field (weak). Number below each level is the degeneracy
of level, number above is the energy.
The different energy levels of a dn configuration are given in Table
3.3 For the complete analysis of the spectra, a correlation diagram
showing all.
Table 3.3 The Ground State and Higher Energy Terms for the dnIon
Configuration
d1, d2
d2, d8
d3, d7
Ground State
Term
2
D
3
F
4
F
d4, d6
5
d5
6
D
S
91
Higher Energy Terms
3
P, 1G, 1D, 1S
4
P, 2H, 2G, 3F, 2 x 2D, 2P
3
H, 3G, 2 x 3F, 3D, 2 x 3P, 1I, 2 x 1G,
1
F, 2x 1D, 2x 1S
4
G, 4F, 4D, 4P, 2I, 2H, 2 x 2G, 2 x 2F,
3 x 2D, 2P, 2S
the energy terms (ground state as well as the excited state) should
be shown (Fig. 3.4 for a d2 ion) and analyzed.
Fig. 3.4 Correlation diagram for d2 system in octahedral field. (a) free
energy terms, (b) terms in a weak field, (c) variations in the field of
intermediate strength, (d) terms in a strong field configurations.
Before we use these correlation diagrams or their simplifications,
knowledge of the selection rules is necessary, as these control the
electronic transition.
3.3.2 Selection Rules
There are two selection rules for electronic transitions in
complexes; these are similar to the selection rules in atomic spectroscopy.
They are :
(a)
Spin Multiplicity Rule : Transitions between states of
different multicity, S, are forbidden. Usually this means that the number
92
of unpaired electrons must not be changed, but this is not quite the same
thing. For example, the transition S2  S1P1 is spin-allowed so long as
the spins of the two electrons in the S1P1 state are + ½ and – ½ (i.e. the
state is a singlet); transition to the triplet state, in which both spins have
the same sign, is forbidden.
Breakdown of the Selection Rules: Spin Multiplicity Rule
The spin forbidden transitions are observed even for the d5 ion
complexes having high pairing energies, but their intensities are very low.
The spin forbidden transitions take place due to the spin-orbit angular
momentum coupling that changes the energies of the different states. Due
to the slight mixing (even 1 per cent) of two states, say a singlet and a
triplet state, the ground state becomes 99 per cent singlet and 1 per cent
triplet, and the excited state becomes 99 per cent triplet and 1 per cent
singlet. The band intensity is then derived from the singlet-singlet and the
triplet-triplet transitions. The extent of mixing depends upon the
differences in their energies and the spin-orbit coupling constants. Thus
the octahedral spin free complexes of d5 ions (Mn2+, Fe3+) must gain
whatever intensity they can through the breakdown of the spin
multiplicity rule as all the higher excited states have lower spin
multiplicity than the ground state 6S.
Laporte Selection Rule :
Transitions within a given set of p or d orbitals (i.e. transitions
involving only redistribution of electrons in the given sub-shell) are
forbidden if the molecule or ion has a centre of symmetry. A more formal
statement of this rule, first put forward by Laporte, is that in a molecule
which has a centre of symmetry, transitions between two g states or
between two u states are forbidden (g = symmetrical; u = unsymmetrical).
93
Breakdown of Laporte's Selection Rule: The mixing of the
orbitals on the metal ion (d-f, d-p), or vibronic coupling can result in the
breakdown of the Laporte's selection rule. Thus if d and p orbitals mix.
 =  (3d) +   (4p),
Where  is the coefficient of mixing. They can mix by producing a
temporary distorted field due to the ligand vibrations, so the atom M is
not at the centre of the symmetric field all the time during which the
electron transition takes place (Frank – Condon Principle).
Hence, in the tetrahedral complexes with no inversion centre, d-p
mixing leads to more intense absorption bands than those for the
octahedral complexes. The absorption of [MX4]2- where M = Ni, Co, Cu,
and X = Cl, Br, I is 20-120 1 mm-1 mol-1 whereas in the octahedral
complexes, these Laporte forbidden transitions have absorption of ca 5
units only.
The vibrational and electronic coupling, called the vibronic
coupling removes the centre of symmetry. If a forbidden band lies near an
allowed band (due to permitted transition), the mixing of the fully
allowed and forbidden energy states due to the vibronic coupling gives
intense absorptions. This is called intensity stealing and depends upon the
energy differences between the two states.
94
Check your Progress-1
Notes:(i)
Write your answers in the space given below.
(ii)
Compare your answers with those given at the end of the
unit.
(a)
The ground state Terms for the following free ions and their
splitting components in an octahedral field are :
(b)
dn ion
Ground Term
Splitting Components
d1
-------------
-------------
d3
-------------
-------------
d5
-------------
-------------
With what the two selection rules are related? Name the
complex in which the rule is violated.
Rule
Related with
Breakdown
1st Rule
-------------
-------------
-------------
-------------
-------------
-------------
-------------
-------------
2nd Rule
3.4
ORGEL DIAGRAMS
Orgel diagrams are popular methods to represent the ground - and
the excited – states of a configuration. Similar to the correlation
diagrams, in orgal diagrams also energies of states are plotted against the
field strengths. Although Orgel diagrams are very simple, but as in these
diagrams excited states of different multiplicity, other than the ground
state, are not expressed, these diagrams are applicable to weak field cases
only.
95
If one plots the magnitude of splitting of the energy levels with the
increasing ligand field for a dn system, and take into consideration the
spin-orbit coupling and mixing of the different energy states, especially
under a strong field, the Orgel diagram for the ion is obtained. These are
given for the dn ions in Fig. 3.5
Fig. 3.5 Orgel diagram for the P and F states in
octahedral as well as in the tetrahedral field.
Further, the energy level diagram for the d n ion in the tetrahedral
field has the same form as that for the d10-n system in the octahedral field,
except that the tetrahedral field splitting is considerably less (44 per cent
only). When the octahedral and tetrahedral field splittings are plotted in
the same diagram, the Orgel diagram for the system results (Fig. 3.5).
The fig. 3.5 represents Orgel diagram for Co2+(d7) in the tetrahedral
and octahedral fields. Here again we can see the inverse relationship
between these two symmetries. This is because the tetrahedral field, in
fact is a negative octahedral field. In this diagram effects of the mixing of
the terms can be seen.
It is a common rule, only the terms of the same symmetry are
mixed; while the limit of mixing is inversely proportional to the
96
difference in the energy between these terms. For Co2+ these terms are of
two, 4T1 (tetrahedral) and 4T1g (octahedral) levels. The mixing of terms
takes place, just similar to the mixing of molecular orbitals. The energy of
the higher level is incrased, while that of the lower level decreases. In the
fig. (for CO2+) this is represented for the pairs of 4T1g and 4T1 by
diverging lines, in case the mixing does not take place, the state is
represented by dotted lines. This can be seen, in case of tetrahedral, the
absence of mixing brings the energies of two 4T1 terms gradually closer
with the increasing field strength. While in octahedral complexes just
reverse of this is seen. Thus, in tetrahedral complexes the limit of mixing
is high.
Orgel diagrams are simple means to assess the number of spin
allowed absorption – bands for a complex in an ultraviolet or visible
spectrum.
3.5
TANABE SUGANO DIAGRAMS
In the strong field cases, the term values for the free ions fail to
give the field splitting. Rather, the attempts are made to estimate the
Racah parameters by placing electrons one by one in the t 2 and eg orbitals
and this proves to be a difficult job.
The Tanabe-Sugano diagrams are the plots of the energies in terms
of E/B (B-Racah parameter) against Dq/B, the ground state of the system
being always plotted as abscissa. The diagrams are drawn to a specific
C/B ratio and are therefore not valid for all the similar systems which
generally differ in C/B ratios. For the d4, d5, d6 and d7 systems, the change
of the ground state from the high spin complexes to the low spin
complexes is shown by a vertical line (Fig. 3.6).
97
Dq/B (a)
Dq/B (b)
Fig. 3.6 Tanabe-Sugano diagrams for (a) d2 and (b) d5 ions in
octahedral field.
The Tanabe-Sugano diagrams suffer from the disadvantage that (i)
the diagram depends on C/B ratio; (ii) There is no accurate way to
determine C and B values for the metal ions in complexes.
98
Even though the C and B values for the metal ions in complexes
are lower than those in the free metal ions, the Orgel as well as the
Tanabe Sugano diagrams do not consider this points at all.
Bond Widths and Shapes
In general, room temperature absorption bands are about 1000 cm-1
wide for the d-d transitions due to (i) molecular vibrations. (2) spin-orbit
coupling, and (3) Jahn-Teller distortions. However, if the symmetric field
has already been reduced to field of low symmetry, Jahn-Teller
distortions may be unnecessary, e.g. for different ligands.
The d-d transitions have low intensities. When they are both spin
forbidden and Laporte forbidden,  is very low. Substitution in
octahedral complexes, which destroys the centre of symmetry of the
ligand, gives higher absorption. Laporte restriction does not apply to the
tetrahedral complexes due to the absence of the centre of symmetry in the
ligand field. Hence, the tetrahedral complexes of the same ions are more
intensely colored than the octahedral complexes. e.g. [CoCl 4]2- is blue
whereas [Co(NH3)6]2+ pink. Multiplicity forbidden transition bands are
sharp, whereas the multiplicity allowed t2geg transition give a broad
band as the excited state has longer bond lengths and the electronic
transitions are faster than the vibrational transitions.
Tanabe Sugano diagram in a simple form, is expressed in Fig. 3.7
for d6 ion in an octahedral stereochemistry. In these diagrams the ground
state is taken as the abscissa and other energy states are recorded relative
to this abscissa. While the interelectronic repulsion is expressed in terms
of Racah parameters, B and C. The constant B is sufficient to give
information about the energy differences in the levels of same spinmultiplicity, but for the terms of different multiplicity, both B and C
parameters become necessary. For the energy difference between the
99
ground state F term of free ion having same spin-multiplicity and the
excited P term is 15 B (As is seen in case of d 2, d3, d7 and d8
configurations).
In Tanabe-Sugano diagrams the energy, E and the field strength, B
is expressed in terms of E/B and  /B respectively. In Fig. 3.7 C/B = 4.8;
while, in the ions of most of the transition-series B is nearly equal to 1000
cm-1 and C = -4B.
Fig. 3.7 Tanabe-Sugano diagrams for d6 ions
The ground state of a d5 ion is 5D, which splits into high energy 5Eg
and low energy 5T2g levels. The low spin state for d6 ion is 1I, which is a
high energy term in the ground state. But as can be seen from the energy
level splitting diagram for the d6 ion, the 1A1g and 1T1g states get
stabilized more than the 5T2g state so that at higher field strength, the
ground state of the ion becomes 1A1g. The spectra of the cobalt (III), a d6
ion can now be predicted easily. The high spin complex [CoF6]3- shows
only one absorption band at 13000 cm-1 corresponding to 5T2g6Eg
transition, whereas the low spin cobalt (III) complexes show two peaks
100
corresponding to 1A1g1T1g and 1A1g1T2g as shown by both [Co(en)3]3+
(21400, 29500 cm-1) and [Co(Ox)3]3- (16500, 23800 cm-1) ions.
3.6
CALCULATIONS OF Dq, B AND  PARAMETERS.
The value of  and B for a given complex may be calculated by
fitting its observed spectrum into the Tanabe-Sugano diagram. In
comparison to the free ion (Bo), the value of electronic-repulsion constant
in its complex (B) is always less. This is due to nephelauxetic effect. The
ratio of B and B0 is known as Nephelauxetic constant  , and is the
measure of covalent character in the complex:

Value of Bin the Complex
B

Bo
Value of Bin the freion
Lesser is the value of  , higher is covalent character in the
complex.
For example, the low spin complex of Co(III) (d6ion), [Co(en)3]3+,
gives two bands at 21550 cm-1 and 29600 cm-1 in its spectrum (Fig. 3.8);
which may be assign to:
21550 cm-1 = 1A1g1T1g transition;
and
29600 cm-1 = 1A1g1T2g transition.
The ratio of these energies will be:1
A1g 1T2 g
1
A1g 1T1g
29600cm 1

 1.37
21550cm 1
101
Fig. 3.8 Electronic spectra of [Co(en)3]3+ and [Co(Ox)3]3- Complexes.
On fitting this into the Tanabe-Sugano diagram (Fig. 3.7), this ratio is
obtained at  / B = 40. From the diagram, the value of E/B obtained for
the transition of lowest energy, is 38. Hence,
1
A1g 1T1g
B

21550cm1
 38
B
This gives, B = 570 Cm-1, which very much less, than the value of
Bo, for the free Co3+ ion (= 1100 cm-1).
From the value of  / B (= 40) and B(=570 cm-1) , the value of 
obtained is 23000 cm-1 (Theoretical value, 23160 cm-1).
102
Check Your Progress-2
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
units.
(a) (i) In Orgel diagrams...................................................are plotted
against..............................................; while in Tanabe-Sugano
diagrams.............................is plotted against...................................
(ii) Orgel diagrams are suitable for complexes of.........................
fields only, but Tanabe-Sugano diagrams are suitable for.............
B and C Racah parameters represent.....................................while,
 represents.......................and is the ratio of .........................
3.7
CHARGE TRANSFER SPECTRA
All electronic transition between orbitals that are centered on
different atoms is called a charge-transfer transition and the absorption
band is usually very strong.
Two types of the charge transfer (CT) bands appear in the metal
complexes:
(1)
The LM CT bands as shown by oxide, chloride bromide and
iodide complexes. For d0 ions, where the d-d transitions are not
possible, the CT bands are used for the determination of the
10Dq values. Thus the three CT bands at 1.85, 3.22 and 4.44 x
104 cm-1 in MnO4- ion, assigned to the t1    e  * , t2   
e  * and t1    t2  *,  * transitions, give the 10Dq for the
oxide ion as e  *  t2  * as 2.59 x104 cm-1 (312 KJ mol-1).
(2)
The ML CT bands that occur in the  acceptor ligands (CO,
NO, CN-) containing empty low energy orbitals.
As the CT bands are neither multiplicity forbidden, nor Laporte
forbidden, they have high absorption intensities (50-2001 mm-1 mol-1).
103
Metal-to-Ligand Charge Transfer:
The visible absorption spectra of iron(II) complexes with ligands
containing the –diimine unit
\
–N
/
N–
have intense charge-transfer bands associated with the transfer of charge
from metal t2g orbitals to the antibonding orbitals of the  –diimine group.
In the case of the 1, 10-phenanthroline complex Fe(phen)32+, the
transition occurs at 19,600 cm-1. A
series of 6-coordinate iron (II)
complexes containing the tertraimine macrocyclic ligand (TIM) shown
below, have been prepared with various monodentate ligands occupying
the two axial sites. The band maxima of these complexes, corresponding
to metal-to-TIM charge transfer indicated that the transition energy is a
function of the  acceptor ability of the axial ligands. As the  -acceptor
ability increases, the energy of the dxz and dyz orbitals decreases relative
to that to the  * (TIM) orbital, causing the t2g  * (TIM) charge-transfer
band to move to higher energies.
Ligand-to –Metal Charge Transfer:
Such bands are often prominent in the spectra of complexes in
which there are electrons in π -orbitals of the ligands. The specra of
RuCl62- and IrBr62- (d4 and d5 complexes, respectively) show two sets of
bands that have been assigned to transitions from the weakly bonding  orbitals on the ligands to the anti-bonding t2g* and eg* orbitals of the
metal atom. In IrBr63- (a d6 complex), the t2g* orbitals are field, and only
the transitions to the eg* orbitals can be observed.
Similarly, ammine and halo (except fluoro) complexes of Co(II)
and Cr(III) give charge-transfer bands at 250 mm or at the higher wave
104
lengths. These indicate LM electron transfer, because the difference of
energy between the lowest energy empty molecular orbitals in the metal
ion and the highest filled molecular orbitlas in the ligand is very less
(lesser than 10,000 cm-1).
For charge transfer bands the values of molar extinction
coefficient,  , are much higher ( x 500-2000) than that of d-d
transitions( =100).
3.8
ABSOLUTE CONFIGURATION IN OPTICALLY ACTIVE
METAL CHELATES
The determination of the absolute spatial relationship (the
chiralitys or "handedness") of the atoms in a dissymmetric coordination
compound is a problem that has intrigued inorganic chemists from the
days of Werner. The latter had none of the physical methods now
available for such determinations. Note that it is not possible to assign the
absolute configuration simply on the basis of the direction of rotation of
the plane of polarized light, although we shall see that, through analysis
of the rotatory properties of enantiomers, strong clues can be provided as
to the configuration.
There are two phenomena associated with these d-d transitions
that are useful in assigning absolute configurations. The two optical
rotatory dispersion (ORD) and circular dichroism (CD), from the basis
for the cotton effect. A general rule may be stated: If in analogous
compounds corresponding electronic transitions shown Cotton effects of
the same sign. the compounds have the same optical configuration.
3.8.1 Optically Rotary Dispersion (ORD):
The effect of a given concentration of a particular optically
active species on the plane of polarization of light is not constant, but
depends upon the wavelength of the light, the direction as well as the
105
extent of rotation being affected. This variation in sign and magnitude
of rotation with wavelength, which is illustrated in Fig. 3.9, is known as
optical rotatory dispersion. If the rotation rises to a maximum towards
short wavelengths before changing sign, the compound is said to show a
positive cotton effect; the opposite behavior constitutes a negative
cotton effect.
As a matter fact, the rotation of the plane polarized light is due to
different refractive indices n1 and nr, for the left and right circularly
polarized light. Hence, the two components interact differently with the
medium and emerge out of phase. On combination, the plane of
polarized light gets rotated. The variation of the rotation of the plane of
plane polarized light with wavelength (Fig. 3.9) is called the optical
rotatory dispersion, (ORD) and the abrupt reversal of the rotation in the
vicinity of the absorption band is called the Cotton effect.
Fig. 3.9 (a) Positive Cotton effect.
(b) Negative Cotton effect.
ORD curves are useful in the assignment of absolute configuration.
For
example,
the
configurations
of
the
enantiomers
of
tris
(ethylenediamine) cobalt (III), and bis (ethylenediamine) glutamatocobalt
(III) are known from X-ray investigations, and it is found that the three
A-(D)-configuration of the ORD spectra to one of known configuration.
For the ORD spectrum to be unambiguous, no other absorption must be
nearby.
106
3.8.2 Circular Dichroism, CD
A chiral molecule displays circular dichroism. That is, the
molecule has different absorption coefficients of right and left circularly
polarized light at any given wavelengths. A Circular dichroism spectrum
(a CD spectrum) is a plot of the difference of the molar absorption
coefficients for right and left circularly polarized light against
wavelength. As we see in Fig. 3.10 enantiomers have CD spectra that are
mirror-images of each other. The usefulness' of CD spectra can be
appreciated by comparing the CD spectra of two enantiomers of
[Co(en)3]3+ (Fig. 3.10 (b)] with their conventional absorption spectra (Fig.
3.10(a). The latter show two ligand field bands, just as thought the
complex were an octahedral complex, and give no sign that we are
dealing with a complex of lower (D3) symmetry. That difference,
however, shows up in the CD spectra where an additional band is seen
near 24000 cm-1.
Fig. 3.10 (a) The absorption spectrum of [Co(en)3]3+ and (b) the CD
spectra of the two optical isomers.
Although ORD was used extensively at one time because of
simpler instrumentation, circular dichroism is currently much more
useful. The CD effect because there is differential absorption of left and
right circularly polarized light associated with transitions such as 1A11E
and 1A11A2. The circular dichroism is the difference between the molar
absorptivities of the left and right polarized 1 r , (solid curves in Fig.
3.9). Complexes having the same sign of CD for a given absorption band
will have the same absolute configuration. Hence, CD spectra can be used
107
to relate large families of chiral complexes to the small number of
primary cases, including [Co(en)3]3+, for which absolute configurations
have been established by X-ray diffraction.
Conformation of Chelate Rings
In addition to the dissymmetry generated by the tris (chelate)
structure of octahedral complexes, it is possible to have dissymmetry ion
the ligand as well. For example, the gauche conformation of
ethylenediamine is dissymetric and could be resolved were it not for the
almost complete absence of an energy barrier preventing recemization.
Attachment of the chelate ligand to a metal retains the chirality of the
gauche form, but the two enantiomers can still interconvert through a
planar conformation at a very low energy, similar to the interconversions
of organic ring systems. Thus, although it is possible in principle to
describe two enantiomers of a complex such as [Co(NH3)4(en)]3+, in
practice it proves to be impossible to isolate them because of the rapid
interconversion of the ring conformers.
If two or more rings are present it one complex, they can interact
with each other and certain conformations might be expected to be
stabilized as a result of possible reductions in interatomic repulsions. For
example, consider a square planar complex containing, two chelated rings
of ethylenediamine. From a purely statistical point of view we might
expect to find three structures, which may be formulated M  , M  and
M  (which is identical to M   ). The first two molecules lack a plane
of symmetry, but M  is a meso form. Corey and Bailar were the first to
show that the M   and M   should predominate over the meso form.
Check Your Progress- 3
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the unit.
108
(a) (i) LM charge transfer bands arise due to transfer of charge
from..............................to................................and are exemplified
by.............................complexes.
(ii) ML charge transfer are seen in............................ligands
containing...................................e.g........................complexes.
These
bands
arise
due
to
transfer
of
charge
from......................orbitals to ........................orbitals of the ligand.
(b)(i) ORD is the variation in.................................................................
..............with.......................;while
CD
plots..........................
against wave length.
(ii) The absolute configuration in optically active metal chelates may
be assigned using a general rule:.....................................................
.........................................................................................................
........................................................................................................
3.8
LET US SUM UP:
After going through this unit, you would have achieved the
objectives stated earlier in this unit. Let us recall what we have discussed
so far.

Most of the transition metal complexes are highly colored as they
absorb radiant energy in the visible region of the spectrum.

Transition metal spectra are of four types:
(i) d-d spectra: Which occur in the near IR, visible and UV regions,
and are due to excitation of electrons from the t2g level to e*g
levels largely in the metal itself.
109
(ii) LM charge transfer spectra: Takes place from a MO located
primarily on the ligand to a non-bonding or anti-bonding MO
located primarily on the metal atom.
(iii) ML charge transfer spectra: Involve transitions of electron
from an anti-bonding or non-bonding orbitals, concentrated on
the metal atom, to the anti-bonding orbital located primarily on
the ligands.
(iv) Intraligand transitions: Involve electron transitions from one
ligand orbitals to another ligand orbitral.

In the absence of external field d orbitals of the metal ion are
degenerate and their ground states for dn configuration are:

d1
2
d6
5
d2
3
F
d7
4
F
d3
4
F
d8
3
F
d4
5
D
d9
2
d5
6
d10
1
D
S
D
D
S
In presence of the external field the degenerate d-orbitals split into
groups of orbitals according to the field strengths. Orgel and
Tanabe-Sugano diagrams represent the ground and excited states of
a given configuration.

Orgel diagram plot the magnitude of splitting of the energy levels
with the increasing ligand for dn system, and take into
consideration the spin orbit coupling and mixing of the different
energy states. It plots the energy, E, against the field strength, B.
Orgel diagrams fail to explain the spectra fully, especially in cases
of strong fields.

Tanabe-Sugano diagrams are most suitable for explaining spectra
both in cases of weak and strong field. They plot the energies in
110
terms of E/B (B=Racah parameter) against Dq/B, the ground state
of the system being always plotted as abscissa. Thus, they also take
into account of interelectronic repulsion (in the form of B and C
parameters).

The value of  and B for a given complex may be calculated by
fitting its observed spectrum into the Tanabe-Sugano diagrams.
The value of Nephelauxetic coefficient (interelectronic repulsion),
 is obtained by dividing electronic repulsion constant of the
complex, B with that of the free ion, Bo:

B
Bo
From the Tanabe-Sugano diagrams, the value of E/B for the
lowest energy transition may be read (=x). This will give the value
of B:
Lowest energy transition
X
B
From the values of  /B (=X) and B, the value of  is obtained.

There are two phenomena associated with the d-d transition that
are useful in assigning absolute configuration in optically active
metal chelates, Optical Rotatory Dispersion (ORD) and Circular
Dichroism (CD).

The variation of the rotation of the plane of plane polarised light
with wave length is called ORD. While CD spectrum is a plot of
the difference of the molar absorption coefficients for right and left
circularly polarized light against wave length.

ORD and CD form the basis for cotton effect (The abrupt reversal
of the rotation in the vicinity of the absorption band), and as a rule
if in analogous compounds, corresponding electronic transitions
111
show cotton effects of same sign the compounds have the same
optical configuration.
3.10 CHECK YOUR PROGRESS : THE KEY
1(a)
dn ion
Ground Term
d1
2
d3
d5
Splitting Components
d
Eg and T2g
4
F
A2g, T2g and T1g
6
S
A1g (No Splitting)
(b)
1st rule
Spin multiplicity Spin free oh complexes of Mn2+(d5)
2nd rule Redistribution of NiCl42-, Th complex gives more
electrons in a
intense absorption than of an Oh
given sub-shell
complex.
2(a)(i)In Orgel diagrams energy state E are plotted against field
strength. B; while in Tanabe-Sugano diagrams E/B is plotted
against  /B.
(ii) Orgel diagrams are suitable for complexes of weak fields only;
but tanabe-sugano diagrams are suitable for both weak and strong
fields.
(b) B and C Racah parameters of complexes represent inter
electronic repulsion; while nephelauxetic coefficient  represent
nephelauxetic effect and is the ratio of B and B0 i.e.  
B
Bo
3(a)(i)LM charge transfer bands arise due to transfer of charge from
the weakly bonding  orbitals on ligands to the anti-bonding t*2g
and e*g orbitals of the metal atom; and are examplified by
chloride, bromide or iodide complexes.
112
(ii) ML charge transfer is seen in  acceptor ligands containing
empty low energy orbitals, e.g. CO, CN-, NO complexes. These
bands arise due to transfer of charge from metal t2g orbitals to
anti-bonding orbitals of the ligand.
(b)(i) ORD is the variation in sign and magnitude of rotation of plane
of polarized light with wave length; while CD plots the difference
of the molar absorption coefficients for right and left circularly
polarized light against wave length.
(ii) The absolute configuration in optically active metal chelates may
be assigned using a general rule: if in analogous compounds,
corresponding electronic transition shows cotton effects of same
sign, the compounds have the same optical configuration.
113
M.Sc. (Previous) Chemistry
Paper – I :
INORGANIC CHEMISTRY
BLOCK – II
UNIT – 4 : Magnetic properties of Transition Metal
Complexes.
UNIT – 5 : Metal π Complexes
UNIT – 6 : Reaction Mechanism of Transition Metal
Complexes-1
Author – Dr. Purushottam B. Chakrawarti
Edtor – Dr. M.P. Agnihotri
114
UNIT-4
MAGNETIC PROPERTIES OF
TRANSITION METAL COMPLEXE
Structure
4.0
Introduction
4.1
Objectives
4.2
Magnetic Moments
4.3
4.2.1
Number of Unpaired Electrons
4.2.2
Spin Only Formula
Anomalous Magnetic Moments
4.3.1
Orbital Contribution in Magnetic Moments
4.3.2
Curie's Law
4.4
Magnetic Exchange Coupling
4.5
Let Us Sum Up
4.6
Check Your Progress: The Key
115
4.0
INTRODUCTION
Substances were first classified as diamagnetic or paramagnetic by
M.Faraday (1845). But it was not untill many years later that these
phenomenon came to be understand in terms of electronic structures.
When any substance is placed in an external magnetic field, there is an
induced circulation of electrons producing a net magnetic moment
aligned in opposition to the applied field. This is the diamagnetic effect
and it arises from paired electrons within a sample. Paramagnetism is
produced by unpaired electrons in a sample. The spin and Orbital motion
of these electrons give rise to permanent molecular moments that tend to
alignt themselves with an applied field.
Magnetic properties and electronic spectra are closely connected.
Magnetic susceptibility measurements are used to decide between
different electronic configurations. It may be mentioned, although the
electronic spectra is a powerful method for investigating transition metal
complexes, additional and complementary information can be provided
by magnetic measurements.
In this unit we shall discuss how net magnetic moments of
transition metal complexes can be worked out; and in what conditions
anomalous magnetic moments are obtained.
However, it will be advantageous if you recall what you have
already studied earlier about the basic concepts of magnetic moments of
atoms.
116
4.1
OBJECTIVES
The main aim of this unit is to study magnetic properties of transition
metal complexes and to establish their correlation with their spectral
properties. After going through this unit you should be able to:
 calculate magnetic moments and number of unpaired electrons in a
transition metal complex;
 describe under what conditions spin-only formula will be useful to
calculate µ of the complexes;
 discuss under which conditions orbital contributions will be
important to calculate µ of the complexes; and
 explain magnetic exchange coupling and spin crossover to describe
anomalous magnetic moments of some complexes.
4.2
MAGNTIC MOMENTS
When a substance is subjected to a magnetic field, H, a
magnetization, I, is induced. The ratio I/H is called the "volume
susceptibility",
K,
and can be measured by a variety of techniques,
including the Gouy balance method, the Faraday method, and an nmr
method. The volume susceptibility is simply related to the "gram
susceptibility," x, and the "molar susceptibility", xm
x
xM 
K
d
K
M
d
where d and M are the density and molecular weight of the
substance, respectively. For most substances, K, x and xM have negative
values; such substances are weakly repelled by a magnetic field and are
called "diamagnetic". For substances having unpaired electrons that do
not strongly interact with one another, K, x and xM have relatively large
117
positive values; these substances are attracted into a magnetic field and
are called 'paramagnetic."
When a paramagnetic substance is placed in a magnetic field, the
moments of the paramagnetic molecules or ions tend to align with the
field; however, thermal agitation tends to randomize the orientations of
the individual moments. Theoretical analysis of the situation leads to the
relations;
N 2
3kT
x Mcorr 
where X Mcorr is the molar susceptibility which has been corrected
both for the diamagnetic contribution to the susceptibility (due to the non
paramagnetic atoms in the sample) and for any small temperatureindependent paramagnetism arising from paramagnetic excited states of
the system. N is Avogadro's number, k is the Boltzmann constant, µ is the
"magnetic moment" of the molecule, and T is the absolute temperature.
By substituting numerical values for N and k, we obtain;
X Mcorr 
0.125 2
T
or   2.83 X Mcorr T
4.2.1 Number Of Unpaired Electrons
Once an experimental value of xM has been obtained for a
paramagnetic substance, it can be used to determine how many unpaired
electrons there are per-molecule or ion. In order to translate the
experimental result into the number of unpaired spins, it must be
recognized that a measured susceptibility will include contributions from
both paramagnetism and diamagnetism in the sample. Even though the
latter will be small, it is not always valid to consider it negligible. The
most common procedure is to correct a measured susceptibility for the
diamagnetic contribution. Compilations of data from susceptibility
118
measurements on a number of diamagnetic materials make it possible to
estimate the appropriate correction factors. The diamagnetic susceptibility
for a particular substance can be obtained as a sum of contributions from
its constituent unit: atoms, ions, bonds, etc. The basic assumption
underlying such a procedure, namely, that the diamagnetism associatiated
with an individual atom or other unit in independent of environment, has
been shown to be valid.
The next step is to connect the macroscopic susceptibility to
individual molecular moment and finally to the number of unpaired
electrons. From classical theory, the corrected or paramagnetic molar
susceptibility is related to the permanent paramagnetic moment of a
molecule µ, by:
xM 
N 2 2
3RT
where N is Avogadro's number, R is the ideal gas constant, T is the
absolute temperature, and µ, is expressed in Bhor magnetrons (BM) (1
BM = eh/4  m). Solving this expression for the magnetic moment gives:
 3RTX m 
 

2
 N

1/ 2
 2.84( x M T )1 / 2
As we know, this paramagnetic moment in the spins and orbital
motions of the unpaired electrons in the substance. There are three
possible modes of coupling between these components spin-spin, orbitalorbital, and spin-orbital. For some complexes, particularly those of the
lanthanides, we must consider all three types of coupling. The theoretical
paramagnetic moment for such a complex is given by:
  g[ J ( J  1)]1 / 2
119
where J is the total angular momentum quantum number and g is
the Lande splitting factor for the electron, defined as:
g 1
J ( J  1)  S ( S  1)  L( L  1)
2 J ( J  1)
The value of J depends on the total orbital angular momentum
quantum number, L, and total spin angular momentum quantum number.
4.2.2 Spin Only Formula
For complexes in which spin-orbit coupling is nonexistent or
negligible but spin and orbital contributions are both significant, the
predicted expression for µ is;
  [4S (S  1)  L( L  1)]1 / 2
This equation describes a condition that is never fully realized in
complexes because the actual orbital contribution is always somewhat
less than the ideal value. This occurs because the orbital angular
momentum is reduced from what it would be in the free metal ion by
presence of ligands. In the extreme case, where general situation in
complexes having A or E ground states, which would include octahedral
d3, d4 (high spin), d6 (low spin), d7 (low spin) and d8 cases. Furthermore,
when a complex involves a first-row transition element, even if the
ground state is T, the orbital contribution generally may be ignored. For
the L=O condition, the above Eq. reduces to;
  [4S (S  1)1 / 2  2[S (S  1)]1 / 2
which is known as the spin only formula for magnetic moment.
By recognizing that S will be related to the number of unpaired electrons
(n) by S = n/2, the expression may be further simplified to;
  [n(n  2)]1 / 2
120
Check Your Progress - 1
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
1. Molar susceptibility xM, is given by the relation;
xM = ..........................................
2. Magnetic moment, µ is given by the relation;
µ = ..........................................
3. The spin only formula is;
µ = ..........................................
4. Magnetic moment, µ and number of un-paired electrons, n, are
related as;
µ = ..........................................
4.3
ANOMALOUS MAGNETIC MOMENTS
Table 4.1 indicates that the values of magnetic moment calculated
using the spin only formula in number of cases differ from the values
obtained from theoretical considerations. This difference is supposed to
be due to two reasons, firstly due to the contribution of orbital magnetic
moment; and secondly due to dependence of magnetic properties on the
experimental temperature (Curie's Law).
121
Table 4.1: Magnetic properties of some complexes of the first-row
transition metals
Central
metal
Ti3+
High spin complexes
Low spin complexes
No. of d
No. of
No. of
µ(expt) µ(calc)b
µ(expt) µ(calc)b
electrons unpaired
unpaired
BM
BM
BM
BM
electrons
electrons
1
1
1.73
1.73
-
V4+
1
1
V3+
2
2
V2+
3
3
Cr3+
3
3
Mn4+
3
3
Cr2+
4
4
Mn3+
4
4
Mn2+
5
5
Fe3+
5
5
Fe2+
6
4
Co3+
6
4
Co2+
7
3
Ni3+
7
3
Ni2+
8
2
Cu2+
9
1
1.681.78
2.752.85
3.80390
3.703.90
3.8-4.0
1.73
-
-
-
2.83
-
-
-
3.88
-
-
-
3.88
-
-
-
3.88
-
-
-
4.754.90
4.905.00
5.656.10
5.706.0
5.105.70
-
4.90
2
2.83
4.90
2
3.203.30
3.18
5.92
1
1.73
5.92
1
1.802.10
2.0-2.5
4.90
0
-
-
4.90
0
-
-
4.305.20
-
3.88
1
1.8
1.73
3.88
1
1.8-2.0
1.73
2.803.50
1.702.20
2.83
-
-
-
1.73
-
-
-
122
2.83
1.73
4.3.1 ORBITAL CONTRIBUTION IN MAGNETIC MOMENTS
In an octahedral ligand field, only the t2g orbitals remain
degenerate and rotationally related. The eg orbitals get separated by 10Dq.
Hence, orbital momentum due to the dx2-y2 orbital electron gets quenched,
and the spin-only formula should apply.
It can be seen that the orbital angular momentum formula should
be important for the high spin d1, d2, d6 and d7 ion complexes and low
spin d4, d5 ions in octahedral field. In the tetrahedral field, the high spin
d3, d4, d8 and d9 ion should have a significant contribution from the orbital
angular momentum. The magnetic moment of [CoCl4]2- (4.4 BM) and that
of [Co(H2O)6]2+ (5.0) confirm the above statements, the orbital moment
contributes for the high spin octahedral, but not for the tetrahedral
complexes.
Even for the other ions, where no orbital moment is expected,
the observed values significantly depart from the spin-only formula
(though the differences are small). This is attributed to the spin-orbitals
interactions which oppose the quenching of the orbital moments by
mixing the orbitals. This explains the generally, lower µeff values
obtained for Cr2+, Cr3+, V3+ and V+ and higher values for spin free Fe2+,
Co2+, Ni2+ and Cu2+ complexes. Greatest deviations occur for the Co2+
and Fe2+ complexes, for which unquenched orbital moments contribute
significantly.
For the 4d and 5d ions, diamagnetism results for even numbered
electrons, and paramagnetism to the extent of one unpaired electron only
is observed for the old numbered electrons, indicating that spin pairing
takes place for these ions as far as possible. This may be due to (i)
reduced enterelectronic repulsions in larger sized ions reducing the
123
electron pairing energies, (ii) higher LF or MO splittings. The µ at room
temperature is generally lower than µs and cannot be used to determine
the unpaired electrons due to (iii) high spin-orbit coupling constants
which align L and S vectors in opposite directions destroying the
paramagnetism.
Further, (iv) the Curie or Curie-Weiss law does not hold, the
variation of µ with L is complex and depends upon the number of the
electrons present.
Some ions like MnO4-, CrO42- and low spin Co3+ complexes
show temperature-independent paramagnetism (TIP) even though they do
not have any unpaired electron. This is due to the spin-orbit coupling of
the ground state to a paramagnetic excited state under the influence of the
magnetic field. The degree of mixing is independent of temperature but
depends on the applied magnetic field, as the excited state is well
separated from the ground state, whose population does not change with
temperature.
4.3.2 CURIE'S LAW
The observed magnetic moments for the metals in t2g ground
state are temperature dependent and usually depart from the µs value due
to probably the t2g electron delocalization and lower symmetry ligand
field components.
Pierre Curie established in 1895 that paramagnetic susceptibility
is inversely proportional to the absolute temperature.
xM = C/T
This expression, which is known as Curie's Law, is actually a
restatement of magnetic moment. The Curie law is obeyed fairly well by
paramagnetic substances that are magnetically dilute, i.e. those in which
124
the paramagnetic centers are well separated from each other by
diamagnetic atoms. In materials that are not magnetically dilute, unpaired
spin on neighboring atom may couple with each other, a phenomenon
referred to as magnetic exchange. Materials that display exchange
behavior can usually be treated with a modification the Curie-Weiss Law;
xM 
C
(T   )
where ø is a constant with units of temperature. If the interacting
magnetic dipoles on neighboring atoms tend to assume a parallel
alignment, the substance is said to be ferromagnetic (Fig. 4.1(b)). If, on
the other hand, the tendency is for an anti-parallel arrangement of the
coupled spins, the substance is anti-ferromagnetic.(Fig. 4.1(c)) In any
material that exhibits magnetic exchange, the tendency towards spin
alignment will complete with the thermal tendency favoring spin
randomness. In all cases, there will be same temperature below which
magnetic exchange dominates, this temperature is called the Curie
temperature (TC) if the type of exchange displayed is ferromagnetic and
the Neel temperature (TN) if it is anti-ferromagnetic. The change in
susceptibility as the temperature is decreased below either TC or TN may
be quite dramatic.
Paramagnetism
Ferromagnetism
Antiferromagnetism
(a)
(b)
(c)
Fig.4.1 Schematic representations of magnetic dipole
arrangements in (a) paramagnetic, (b) ferromagnetic, and (c) antiferromagnetic materials.
Fe(Phen)2(CNS)2 is an example which shows significant variation in
magnetic moment with temperature. (Fig. 4.2)
125
Fig.4.2 The magnetic moment of Fe(phen)2(NCS)2
as a functions of temperature.
4.4
MAGNETIC EXCHANGE COUPLING
As we know, a number of transition metal ions form both high
and low spin complexes, and we have now seen that magnetic
susceptibility allow us to experimentally distinguish one from the other.
Within ligand field theory, these two spin configurations in octahedral
complexes are explained in terms of relative magnitudes of
 and
pairing energy (P): We associate high spin complexes with the condition
   P and low spin complexes with    P . For complexes in which the
energy difference between  and P is relatively small, an intermediate
field situation, it is possible for the two spin states to coexist in
equilibrium with each other. Consider the Fe2+ ion. At the two extremes,
it forms high spin paramagnetic [Fe(H2O)]2+ (S=2) and low spin
diamagnetic [Fe(CN)6]4- (S=0).
Octahedral complexes with 4, 5, 6 or 7 d electrons can be either
high-spin or low-spin, depending on the magnitude of the ligand-field
splitting,  . When the ligand-field splitting has an intermediate value
such that the two states of the complex have similar energies, the two
states can coexist in measurable amounts at equilibrium. Many
"crossover" systems of this type have been studied.
126
4.5
SPIN CROSSOVER
With the change in field strength, change in the magnetic moment
i.e., the change from high-spin to low-spin can be explained in terms of
splitting of electronic states with the field strength, e.g. the TanabeSugano diagram to these d6 complexes show that near the crossover point
between weak and strong field the difference in energy between the spinfree (5T2g) and spin-paired (1A1g) ground states becomes very small (Fig.
4.3) within this region, it is reasonable to expect that both spin state may
be present simultaneously and that the degree to which each is
represented will depend on the temperature (  - P = kT). A complex
illustrating these effects is [Fe(phen)2(NCS)2] (Fig. 4.2). At high
temperature a moment consistent with four unpaired electrons is
observed, but as the temperature is decreased, a sharp drop in magnitude
is observed at 175K where the low-spin form becomes dominant. Usually
spin transitions occur somewhat more gradually than in the case shown
here, and reasons for the abruptness observed for this complex, as well as
some residual paranagnetism seen at low temperature have been
discussed extensively.
5
T2g
1
A1g
E

Fig.4.3 Variation in energies of 5T2g and 1A1g terms with increasing
6
 for d octahedral complexes. At weak field (high spin complexes)
the ground term is 5T2g, while at strong fields (low spin complexes) it
is 1A1g Note that in the region immediately on each side of the spin
crossover point, the energy difference between the two terms is small;
thus high and low spin complexes coexist.
127
In solutions, these systems are fairly straightforward; the change in
magnetic susceptibility with temperature can be interpreted in terms of
the heat of conversion of one isomer to another. However, treatment of
the system as an equilibrium between two spins yields  H=3.85 kcal
mol-1 and  S = 11.4 for the high spin  low spin conversion. On the
other hand, spin crossover in solids is a complex phenomenon because of
cooperative structural changes and changes in the energy separation of
the high-spin and low-spin states with temperature. Thus the magnetism
of Fe(phen)2(NCS)2 change sharply at 174K, as shown in Fig. 4.2
Check Your Progress - 2
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
(A)(i) Orbital Contribution in magnetic moment is important for high
spin............................ions complexes; and low spin ...................
ions in octahedral field.
(ii)
The
greatest
deviation
in
magnetic
moment
occurs
for...............complexes.
(B) Curie's Law state that.....................................................................
............................................................................i.e. Xm =............
(C) In octahedral complexes for dn configurations (n=...................)
the two states (low-spin and high-spin) of complexes can
coexist in measurable amount's at equilibrium at ligand field
splitting has.
(D) The change from high spin to low spin can be explained in terms
of...................................................................................................
(E) The crossover point is reached when the difference in energy
between..........................................................states become very
small.
128
4.6
LET UP SUM UP
 Substance may be diamagnetic or paramagnetic when any
substance is placed in an external magnetic field, there is an
induced circulation of electrons producing a net magnetic moment
in opposition to the applied field. This is the diamagnetic effect and
it arises from paired electrons within a sample.
 Paramagnetism is produced by unpaired electrons in a sample. The
spin and orbital motion of these electrons give rise to permanent
molecular moments that tends to align themselves with an applied
field.
 When a substance is subjected to a magnetic field, H, a
magnetization I, is induced. The ratio of I/H is called volume
susceptibility, k. The volume susceptibility is simply related to the
'gram susceptibility', x and the molar susceptibility, XM asX
K
d
or X M 
K
M
d
where d and M are the density and molecular weight of the substance,
respectively.
 xM is the molar susceptibility which has been corrected both for the
diamagnetic contribution to the susceptibility and for any smll
temperature-independent paramagnetism from paramagnetism
excited states of the system, and may be given as,
X Mcorr 
0.125 2
or   2.83 X Mcorr T
T
129
when values of Avogadro's number, A, Boltzman constant K,
the magnetic moment of the substance, µ and absolute temperature are
substituted.
 Once an experimental value of XM has been obtained for a
paramagnetic substance, it can be used to determine how many
unpaired electrons there are per molecule or ion.
 For complexes in which spin-orbit coupling is nonexistent or
negligible but spin and orbital contributions are both significant µ
is given by;
μ  [4 S(S 1)  L(L 1)]1/ 2
 When a complex involves a first row transition element, even if the
ground state is T, the orbital contribution generally may be
ignored, and we get L=O and µ is given by spin only formula;
μ  [4 S(S 1)1/ 2  2[S(S 1)]1/ 2
 By recognizing that S will be related to the number of unpaired
electrons (n) by S = n/2, the above expression is simplified to;
μ  [n(n 2)]1 / 2
 Magnetic moment calculated using the spin only formula in
number of cases differ from the values obtained from theoretical
considerations. This deviation may due to be either the contribution
of orbital magnetic moment, or due to dependence of magnetic
properties on the experimental temperature, (Curie's Law).
 The orbital angular momentum formula may be important for the
high spin d1, d2, d6 and d7 ion complexes, and low spin d4, d5 ions
in octahedral field.
130
 The observed magnetic moments for the metals in t 2g ground state
are temperature dependent and usually depart from the µs values
due to probably the t2g electron delocalization and lower symmetry
ligand field components.
 The Curie's Law states that paramagnetic susceptibility is inversely
proportional to the absolute temperature;
xM = C/T
 If the interacting magnetic dipoles on neighboring atoms tend to
assume a parallel alignment, the substance is said to ferromagnetic,
and if, on the other hand, the tendency is for an anti-parallel
arrangement of the coupled spins, the substance is antiferromagnetic.
 Fe(phen)2(CNS)2 is an example which shows significant variation
in magnetic moment with temperature. A number of transition
metal ions form both high and low spin complexes. These two spin
configurations in octahedral complexes are explained in terms of
relative magnitudes of  and pairing energy (P). High spin
complexes are formed when  < P and low spin when  > P.
 For complexes in which the energy difference between  and P is
relatively small an intermediate field situation, it is possible for the
two spin sates to co-exist in equilibrium with each other.
 Variation in energies of 5T2g and 1A1g terms with increasing  for
d6 octahedral complexes show, at weak fields (high spin
complexes), the ground term is 5T2g, while at strong fields (low
spin complexes) on each side of the spin crossover point, the
energy difference between the two terms is small; thus high and
low complexes may coexist.
131
4.7
1.
CHECK YOUR PROGRESS: THE KEY
(i)
XM 
KM
d
(ii)  = 2.84 (XMT)½
(iii)  = [4S (S+1) ]½ = 2 (S(S+1)]½
(iv)  = [n (n+2)]½
2.A (i) High spin d1, d2, d6 and d7 ion complexes, and low spin d4 and
d5 ions.
(ii) For Fe2+ and Co2+ complexes.
B.
Curie's Law states that paramagnetic susceptibility is inversely
proportional to the absolute temperature, i.e. XM= C/T
C.
For dn configurations (n = 4, 5, 6, 7) the ligand field splitting
has an intermediate value.
D.
In terms of splitting of electronic states with the fields strength.
E.
Between spin free 5T2g and spin paired 1A1g ground states
become very small.
132
UNIT-5
METAL
π
COMPLEXES
Structure
5.0
Introduction
5.1
Objectives
5.2
Metal Carbonyls
5.3
5.4
5.2.1
Classification
5.2.2
Isolobal Concept
5.2.3
Methods of Preparation and Properties
5.2.4
Structure
5.2.5
Vibrational Spectra
Metal Nitrosyls
5.3.1
Neutral NO and NO- Complexes
5.3.2
Complexes of NO+
5.3.3
Pure Nitrosyl Complexes
5.3.4
Nitrosyl Carbonyl Complexes
5.3.5
Nitrosyl Halide Complexes
5.3.6
Nitroso Cyanide Complexes
Dinitrogen Complexes
5.4.1
5.5
Fixation of Nitrogen
Dioxygen Complexes
5.5.1
Heme Proteins and Transportation of O2
5.5.2
Haemoglobin
5.6
Tertiary Phosphine as Ligand
5.7
Let Us Sum Up
5.8
Check Your Progress: The Key
133
5.0 INTRODUCTION
π-bonding in complexes was proposed for the first time, by Pauling

(1924), in the form of back-bonding (M 
L) to account for electro-
neutrality of metal to ligand bond. According to him, if the ligand, linked
with the metal ion through LM, ϭ-bond, has vacant π-orbitals, it can
accept lone pair of electrons from metal-ion (if present) to form ML, πbonds. This also accounts for the extra stability of metal complexes with
unsaturated ligands. However the latest and the most successful theory of
bonding for metal-complexes Ligand Field Theory (LFT), explained
quantitatively while Mπ-bonding stabilizes, the complex, LM
π-bonding destabilize it. This also explains positions of CN- and Fligands in the spectrochemical series.
Most transition metals form complexes with a wide variety of
unsaturated molecules such as carbon monoxide, nitric oxide, dinitrogen,
dioxygen etc. In many of these, the metal is in zero or another low
oxidation state and, as we have already mentioned, π-bonding between
the metal and the ligands is believed to play an important part in
stabilizing these complexes. In this regard metal carbonyls are important
as they involve both metal carbon ϭ and π-bonds. In this unit we shall
consider the metal carbonyls, anions derived from them, some of their
substitution product, and complexes formed by a few other ligands.
5.1 OBJECTIVES
The main aim of this unit is to study π-complexes of transition
metals, with special reference to bonding and their structures. After going
through this unit you should be able to:
134
 describe metal carbonyls, their classification, methods of
preparation and reactions; with special reference to their structures,
 discuss how these complexes, almost without exception, conform
to the effective atomic number rule and isoloble concept,
 explain bonding in these complexes in terms of IR spectra;
 describe preparation, properties and structures of metal nitrosyls,
 discuss dinitrogen complexes and their importance in the fixation
of nitrogen;
 explain formation of dioxygen complexes with special reference to
transportation of oxygen by heme proteins; and
 describe the nature of complexes with tertiary phosphine as a
ligand.
5.2 METAL CARBONYLS
The compounds formed by the combination of CO molecules with
transition metals are known as metallic carbonyls.
Carbon mono-oxide posses a unique property of unsaturation by
virtue of which it may combine with a large number of metals under
suitable conditions. Such compounds of CO with metals are termed as
metallic carbonyls.
In carbonyls, a metal atom is directly linked to the carbon atom of
a carbonyl group. Since the electrons forming OCM bond are supplied
solely by CO molecule, metal atom in carbonyls is said to be in zero
oxidation state. In metal carbonyls CO molecules act as neutral ligands.
Metal carbonyls vary considerable in their properties ranging from
volatile nonpolar to the nonvolatile electrovalent carbonyls. For examplenickel forms volatile nonpolar carbonyls, where as alkali and alkaline
earth metals from non-volatile electrovalent carbonyls.
135
The general formula of the carbonyls may be given as Mx(CO)y
where M is a metal capable of forming carbonyl. Metal carbonyls may be
regarded as parents of number of related compounds such as metal
nitrosyl carbonyl, M (NO)y (CO)x, and metal carbonyl hydrides HxM
(CO)y.
5.2.1 Classification
Carbonyls are classified into two distinct groups:
a.
Monocular carbonyls: These carbonyls have the general
formula Mx(CO)y which contain more than one metal atom
per molecule.
b.
However the carbonyls having 2 metal atoms are called
binocular carbonyls, and
c.
those having more than two metal atoms as ploynuclear
carbonyls. Polynuclear carbonyls may be homonuclear e.g.
[Fe3(CO)12 or heteronuclear e.g. MnCo(CO)9, MnRe(CO)10]
(Table 5.1)
They have following characteristics:
i. These are almost insoluble in organic solvents.
ii. Many polynuclear carbonyls decompose at or below the
melting point.
5.2.3 Preparation And Properties Of Carbonyls
a. Direct synthesis from metals and carbon mono-oxide, for
example:
1.
Nickel reacts with CO at room temperature and normal
pressure;
C
Ni  4CO 40


 Ni(CO) 4
136
2.
When CO is passed over reduced iron at 108o-220o and
pressure of 50 to 200 atom pressure. Fe(CO)5 is formed;
Fe - + 5CO
Fe(CO)5
Rhenium, osmium and iridium carbonyls could not be prepared
by direct reactions.
Table 5.1 The binary carbonyls
Electrons needed
13
to attain noble
gas
configuration
First transition V(CO)6
series
12
Cr(CO)6
11
10
Mn2(CO)10
9
Fe(CO)5
CO2(CO)8
blue solid white solid yellow solid yellow liquid
(sublimes)
(m.p.154o)
(b.p.103o) Fe2 (m.p.51o)
liquid
(CO)12 black
CO6(CO)16
(b.p.43o)
solid
black solid
Second
Mo(CO)6
Te2(CO)10
Ru(CO)5
transition series
white
white solid
colourless
Rh2(CO)8*
solid(subli
liquid
mes)
22o)
orange solid
Ru2(CO)9*
Rh6(CO)16
Ru3(CO)12
black solid
(m.p.- Rh4(CO)12
orange
solid
(m.p.-154o)
W(CO)6
Re2(CO)10
series
white
white
solid(subli
mes)
Os(CO)5
Ir2(CO)8
solid colourless
yellow solid
o
liquid
Ir4(CO)12
(m.p.15o)
yellow solid
Os2(CO)9
Ir6(CO)16 red
(m.p.177 )
Orange
solid solid
Os2(CO)12
Yellow
solid
o
(m.p.224 )
137
Ni(CO)4
orange solid colourless
(sublimes)
Third transition
8
b. Indirect synthesis involving the Gringed reagent: job prepared
chromium hexacarbonyl by the action of CO on a mixture of
grignard reagent and anhydrous chromium chloride in ether
solution.
According to Hiber the primary reaction is as follows:
C6H5MgBr + CrCl3 + CO  Cr(CO)2(C6H5) + MgBrCl + MgBr2
The unstable intermediate compound is composed with acid
to yield the hexacarbonyl:
3Cr(CO)2 (C6H5)4 + 6H  Cr(CO)6 + 2Cr3+ + 12C6H-5 + 3H2
The reactions gives low yield which can be improved by
using high carbon mono-oxide pressure.
c. Indirect synthesis involving metal compounds: Metal carbonyls
can be prepared by the reaction of CO with certain metal
compounds for example:
i. Nickel tetracarbonyl may be prepared by passing CO into a
suspension of nickel cyanide, sulphide or mercaptide suspended in
NaOH solution.
2NiX4 + 2nCO  2Ni(CO)nX + X2
Ni(CO)nX + (4-2n)CO  Ni(CO)4 + NiX2
ii.Ruthenium pentacarbonyl may be prepared by the action of CO
and Rul3 in the presence of an iodine acceptor:
CO]
CO]
RuI 3 [
 Ru (CO) 4 I 2 [
 Ru (CO)5
Similarly [Ir(CO)4]7 may be prepared.
d. Synthesis by carbonylating the metallic salts with CO in the
presence of reducing agent. When salts like R4I3, CrCl3, VCl3 are
made to treat with CO in presence of a suitable reducing agent like
Mg, Ag, Cu, Na, H2 etc.
138

 Cr(CO)6 + LiCl + AlCl3
CrCl3  CO  LiAlH y 115
o
C 250atmpressure


 2Ru(CO)5 + 6Agl
2Rul3 + 10CO + 6Ag 175
atmpressure
2Mnl3 + 10CO + 2Mg 25
C210

 2Mn2(CO)10+ 2Mgl
e. Synthesis from other carbonyls: when iron pentacarbonyl is
exposed to UV light it loses CO and forms Fe2(CO)9. This
compound undergoes thermal decomposition to yield iron
pentacarbonyl and trimeric tetracarbonyl.
.V .

 Fe2(CO)4 + CO
2Fe(CO)5 U

 Fe(CO)5 + [Fe(CO)4]3 + CO
2Fe2(CO)9 heat
f. Synthesis from Carbonyl hydrides: when iron carbonyl hydride
is oxidised by MnO2 or H2O2,[Fe(CO)4]3 is formed.
g. By treatment of oxide of metals with CO under pressure:
Carbonyls of osmium and rhenium are prepared by the reaction of
CO with their oxides under pressure.
C
100


OsO4 + 9CO
Os(CO)5 + 4CO2
50 atm pressure
C
Re2O7 + 17CO 75


 Re2(CO)10 + 7CO2
atm
200
h. Preparation of Mo (CO)6 and W(CO)6 from Fe(CO)5
MoCl6 + 3Fe(CO)5 
 Mo(CO)6 + 3FeCl2 + 9CO
i. Preparation of Fe2(CO)9 and Os2(CO)9 from Fe(CO)5 and
Os(CO)5-9 with cooled solution of Fe(CO)5 and Os(CO)5 in glacial
CH3COOH is irradiated with u.v. light, Fe(CO)9 and Os2(CO)9 are
obtained respectively.
.V.light
2Fe(CO)5 U

 Fe2(CO)9 + CO
.V .light
2Os(CO)5 U

 Os2(CO)9 + CO
139
Properties of Carbonyls
i. The metal carbonyls are crystalline solids, except for nickel
carbonyl and the pentacarbonyls of iron, ruthenium and osmium
which are liquids.
ii. Many are coloured for example: Crystals of cobalt carbonyl are
orange and iron pentacarbonyls is yellow oil and nicked carbonyl is
colourless.
iii. Due to their covalent nature renders them insoluble in water, most
of them are soluble in solvents like CCl4.
iv. Excepting V(CO)6 all the carbonyls are diamagnetic. V(CO)6 is
paramagnetic and its paramagnetic property corresponds to the
presence of one unpaired electron. The metal in carbonyls are in
zero oxidation state.
Table 5.2 Colour And Melting Points Of Some Carbonyls
Carbonyl
V(CO)6
Melting Point, (oC)
Colour and shape
Black crystals
Decomposes
at
70oC,
Sublime
in
vacuum
Cr(CO)6
Colourless crystals
Sublime in vacuum
Mo(CO)6
Colourless crystals
Sublime in vacuum
W(CO)6
Colourless crystals
Sublime in vacuum
Mn2(CO)10
Golden crystals
154o-155o
Re2(CO)10
Colourless crystals
Sublime at 140o and decompose at 177oC
Fe(CO)5
Yellow Liquid
B.P. 103oC
Fe2(CO)9
Bronze Mica-like
Decomposes at 100oC
platelets
Fe3(CO)12
Dark green crystals
Decomposes at 140oC
CO2(CO)8
Orange crystals
51oC
Ni(CO)4
Colourless Liquid
B.P. 43oC
140
Chemical Properties
1. Substitution Reactions: Some or all CO groups present in
carbonyls can be replaced by monodentate ligands such as alkyl or
aryl isocyanide (CNR) PR3, PCl3, Py, CH3OH etc.
Ni(CO)4 + 4CNR  Ni(CNR)4 + 4CO
Ni(CO)4 + 4PCl3  Ni(PCl3)4 + 4CO
Fe(CO)5 + 2CNR Fe(CO)3 (CNR)2 + 2CO
2. Action of NaOH or Na metal: Formation of carbonylate ion:
Aqueous alcoholic solution of NaOH reacts with Fe(CO)5 to form
carbonylate anion [Fe(CO)4]-.
Fe(CO)5 + 3NaOH  Na+[H+Fe2-(CO)4]-Na2Co3 + H2O
H-atom in [H+Fe2-(CO)4]- ion is acidic which implies that Fe
atom in this ion is in -2 oxidation state.
Na-metal in liquid NH3 is able to convert Fe2(CO)9.
Co2(CO)8, Fe3(CO)12, Cr (CO)6,.Mn2(CO)10 etc, into carbonylate
anions and in this conversion these carbonyls are reduced.
Fe2(CO)9 + 4Na  2Na 2 [Fe2-(CO)4]2- + CO
Co2(CO)8 + 2Na  2Na+[Co-(CO)4]4
3. Action of halogens: Most of the carbonyls react with halogens to
yield carbonyl halides. For example:
Fe(CO)5 + X2  Fe(CO)4X2 + CO
Mo(CO)6 + Cl2  Mo(CO)4Cl2 + 2CO
Mn2(CO)10 + X2(X = Br, I)  2Mn(CO)5X
Both Co2(CO)6 and Ni(CO)4 are decomposed into metallic
halides and CO when treated with halogens.
Co2(CO)8 + 2X2  2CoX2 + 8CO
Ni(CO)4 + Br2  NiBr2 + 4CO
141
4. Action of NO: many carbonyls react with nitric oxide (NO) to
form metal carbonyls nitrosyls. For example:
C
Fe(CO)5 + 2NO 95


 Fe(CO)2(NO)2 + 3CO
3Fe3(CO)9+4NO 
 2Fe(CO)2(NO)2+ Fe(CO)5+Fe3(CO)12+6CO
5. Action of H2: Formation of carbonyl hydrides (reduction): when
Mn2(CO)10 and Co2(CO)8 react with H2, they get reduced to
carbonyl hydrides, Mn(CO)5H and Co(CO)4H respectively.
+ 0
C , 200atmpressure
Mn2(CO)10 + H2 200


 2[Mn (CO)5H ]
+
C , 200atmpressure
Co2(CO)8 + H2 165

 
 2[Co (CO)4H ]
6. Action of heat: Different carbonyls yields different products when
heated for example:
C
Fe(CO)5 250

 Fe + 5CO
C


 3Fe(CO)5 + Fe3(CO)12
3Fe2(CO)9 70
C
Fe3(CO)12 140

 3Fe + 12CO
Metal Carbonyls of Different Groups:
1.
Carbonyls of Sixth B Metals
These form carbonyls of one type only M(CO)6 where M =
Cr,Mo, Or W, but chromium also forms Cr(CO)5+
A. Chromium Hexacarbonyl Cr(CO)6.
Preparation:
i.
It is prepared by job's method by passing CO at 50 atm.
pressure and at room temperature into a suspension of chromic
chloride in ether. Which has been treated with phenyl
magnesium bromide at -70oC.
ii.
Chromium hexacarbonyl can be prepared by treating a solution
of a chromic salt dissolved in ether with Al(C2H5)3 and carbon
mono-oxide at a high temperature and pressure.
142
iii. It may also be prepared by carbonylating CrCl3 with CO in the
presence of a reducing agent like LiAlH4.
C , 200atmpressure
CrCl3 + CO + LiAlH4 175

 
 Cr.(CO)6 + LiCl + AlCl3
Properties:
1.
Chromium hexacarbonyl exists in colourless rhombic crystals
which sublime without decomposition and dissolve in either,
chloroform, CCl4 and benzene.
2.
It is attacked by air, bromine, cold aqueous alkali, dilute acids
conc. HCl and Conc.H2SO4. It is however decomposed by
chlorine or by conc. nitric acid.
3.
Decompositions: It gets decomposed by F2 at -75oC to form
CrF6.
4.
Action of Na-Metal or NaBH4: Cr(CO)6 when is treated with
Na metal or NaBH4 in liq. NH3 carbonylate anion is formed. In
these reactions the carbonyls are reduced.
22Liq.NH

 Na 2 [Cr (CO)5] + CO
Cr(CO)6 + 2Na 
3
22NaBH / Liq. NH
Cr(CO)6 
 
 Na 2 [Cr (CO)10] + 2CO
4
5.
3
Substitution reactions: Some CO groups present in Cr (CO)6 can
be replaced by pyridine to get a number of products.
py

Cr (CO)3 ( Py )3

 Cr(CO)4(Py)2 Py

 Cr2(CO)7(Py)5 
Cr(CO)6 Py
Yellow brown solid Yellow red solid Bright red solid
B.
Molybdenum Hexacarbonyl and Tungsten Hexacarbonyl.
Preparation:
1.
Both these carbonyls may be prepared by job's method which
involve the reaction of either MoCl6 or WCl6 with CO in the
presence of phenyl magnesium bromide.
143
2.
Both may also be prepared by the action of CO at 225 o and 200
atm. pressure on metallic molybdenum or tungsten reduced in
the presence of copper or iron.
Properties:
1.
They are colourless, Mo(CO)6 sublimes at 40oC and boils at
156.4o, whereas W(CO)6 sublimes at 50oC and boils at 175oC.
2.
They are stable in air and dissolve in organic solvents like ether,
chloroform, CCl4 and benzene.
3.
Mo(CO)6 do not react with air, cold aqueous alkali, acids,
except conc. nitric acid or with thiols or nitric oxide.
4.
Bromine and chlorine can decompose Mo(CO)6 and W(CO)6.
5.
With pyridine, phenanthroline and ethylene diamine, the CO
group in Mo(CO)6 and W(CO)6 is replaced.
M(CO)6 Pyridine
 M(CO)5Pyr2  M2(CO)7PYr5  M(CO)3PYr3
2.
Carbonyls or VII Group:
These form volatile carbonyls of the formula M2(CO)12
where M = Mn, Te and Re.
Manganese carbonyl, Mn2(CO)10.
Preparation:
1.
This is prepared by treating manganese iodide and magnesium
with CO in ether under high pressure. In this reaction,
magnesium acts as a reducing agent.
C, 210atmpressure


 Mn2 (CO)10
2MnCl2+10CO+2Mg (in diethyl ehter) 250
+ 2MgI2
2.
By carbonylating MnCl2 with CO in presence of (C8H5)2CONa
C140atm.

 Mn2(CO)10+
2MnCl2+10CO+4(C8H5)2CONa 165
4(C6H5)2CO+4NaCl
144
Properties:
1.
Manganese carbonyl forms volatile, golden yellow, crystalline,
solid which melts at 155oC in a sealed tube. It is soluble in
organic solvents. It is slowly oxidised in air, especially in
solution.
2.
Action of halogens: Mn2(CO)10 reacts with halogens to form
carbonyl halides. Mn2(CO)10 + X2 (X=Br2I)  2Mn(CO)5X
3.
Action of Na-Metal: Na-metal in liquid NH3 converts
Mn2(CO)10 into carbonylate anion. In this reaction the oxidation
state of Mn decreases from zero to -1;
Liq.NH

 2Na  [Mn (CO)5]
Mn2(CO)6 + 2Na 
3
4.
Action of H2: Mn2(CO)10 gives carbonyl hydride, Mn(CO)5H;
in the formation of this compound the oxidation state of Mn
decreases from zero to -1.
+ 0
200C , 200atmpressure
 2[Mn (CO)5H ]
Mn2(CO)10 + H2 
5.
Substitution Reaction: Mn2(CO)10 reacts with PR3 to form
Mn(CO)4(PR3):
Mn2(CO)10 + PR3  2Mn(CO)4(PR3) + 2CO
6. Diamagnetic nature: Mn2(CO)10 is a diamagnetic substance,
diamagnetic character is confirmed by the fact that all the
electrons in Mn2(CO)10 are paired and Mn-Mn bonds is present
in it.
3.
Carbonyls of VIII Group Metals
A. Carbonyls of Iron: Three carbonyls of iron are known, these are:
a. Iron Pentacarbonyl, Fe (CO)5
145
Preparation:
i.
It can be prepared by the action of CO on iron powder at 200 oC
and 200 atm. pressures.
Fe + 5CO  Fe (CO)5
ii.
Recently it has been prepared by the action of CO on Ferrous
iodide in the pressure of Cu which acts as a halogen acceptor.
200C
FeI2 + 4CO 
 Fe (CO)5I2
iii. It may also be prepared by the action of CO on FeS at 200 oC
and 200 atm. pressure in the presence of copper.
200C , 200atmpressure
 2Fe(CO)5 + Cu2S
2FeS + 10CO + 2Cu 
Properties:
i.
Fe(CO)5 is a yellow liquid which is soluble in methyl alcohol,
ether, acetone and C6H6. It is insoluble in H2O.
ii.
Decomposition: M thermal decomposition at 250oC it yields
pure Fe.
250C
 Fe + 5 Co
2Fe(CO)5 
iii. Action of u.v. light: When cooled solution of Fe(CO)5 in
glacial CH3 COOH is irradiated with u.v. light, Fe(CO)9 is
formed. The above reaction is reversed in darkness.
iv. Hydrolysis: Fe(CO)5 gets hydrolysed by H2O and acids
Fe (CO)5 + H2SO4  FeSO4 + 5CO + H2
v.
Action of alkali: Fe(CO)5 + 4NaOH  Na+[Fe2-(CO)4H+]- +
Na2CO3 + H2O
vi. Action of NH3: with NH3 it yields Fe(CO)4H2
Fe(CO)5 + H2O + NH3  Fe (CO)4H2 + NH2COOH
146
vii. Reaction with halogen:
Fe(CO)5 + X2  Fe (CO)4X2 + CO
The velocities of these reactions have been found to follow the
order CI < Br < I.
b.
Iron Enneacarbonyl, Fe2(CO)9
When iron pentacarbonyl is dissolved in glacial acetic acid
and is exposed to u.v. light for 6 hours, Fe2(CO)9 is formed which
dissolves in acetic acid, on cooling with water, golden crystals of the
enneacarbonyl are precipitated and are filtered off.
2Fe(CO)5  Fe2(CO)9 + CO
Properties
i. Fe2(CO)9 forms golden triclinic crystal, it is diamagnetic and
non-volatile. It is insoluble in water but soluble in toluene and
pyridine. When heated to 50oC decomposes to form Fe2(CO)12.
3Fe2(CO)9  3Fe (CO)5 + Fe3(CO)12
ii. Action of heat: when heating is done at 100oC, Fe2(CO)9
decomposes to form iron, CO and some Fe (CO)12.
4Fe2(CO)9  Fe + Fe(CO)5 + Fe3(CO)12 + CO
iii. Action of NO: With NO it gives Fe(CO)2 (NO)2 together with
Fe(CO)5 and Fe2(CO)12.
3Fe2(CO)9+4NO2Fe(CO)2(NO)2+Fe(CO)5 + Fe3(CO)12 + 6CO
c. Iron Dodecarbonyl, Fe3(CO)12.
Preparation
It can be prepared by heating Fe2(CO)9 dissolved in toluene at 70oC.
3Fe2(CO)9  3Fe(CO)5 + Fe3(CO)12
147
Properties
i. Fe3(CO)12 forms deep crystals which are soluble in organic
solvents like toluene, alcohol, ether and pyridine.
ii. Action of Heat: When heated to 140oC Fe3(CO)12 decomposes
to give metallic iron and CO.
C
Fe3(CO)12 140

 3Fe + 12CO
iii. Reaction with Na: Carbylate axion is formed when Fe3(CO)12
reacts with Na metal in Liq. NH3.
+
22
Liq.NH
Fe3(CO)12 + 6Na 

 3Na2 [Fe (CO)4]
3
iv. Substitution Reactions: This reaction takes place with pyridine
and methyl alcohol.
Fe3(CO)12 + 3Py  Fe3(CO)9(Py)3 + 3Fe(CO)5
B.
Carbonyls of Cobalt
It forms two carbonyls
i. Cobalt Octacarbonyl, CO2(CO)8
Preparation
i. It is prepared by the action of CO and the reduced metallic
cobalt at 220oC and 250 atm.
2Co + 8CO Co2(Co)8
ii. When a solution of cobalt carbonyl hydride is treated by an
acid, hydrogen is evolved and Co2(CO)8 remains
2Co(CO)4H  Co2(CO)8 + H2
Properties
1. Cobalt octacarbonyl forms orange transparent crystals.
2. It is insoluble in water but is soluble to some extent in alcohol,
ether, CS, etc.
148
3. Action of air: On exposure to air, dicobalt octacarbonyl is
converted into deep violet basic carbonate of cobalt.
4. Action of Na-metal in liq. NH3: When CO2(CO)8 reacts with
Na-metal in liq. NH3, it gets reduced to carbonylate anion.
Co2(CO)8 + 2Na lig
NH
 2Na[Co(CO)4]
3
5. Action of NO: Co2(CO)8 reacts with NO at 40oC to form
cobalit carbonyl nitrosyl, [Co-(CO)3(NO)]0. Thus in this
reaction the oxidation state of cobalt decreases from 0 to -1.
Co2(CO)8 + 2X2  2Co2X2 + 8CO
6. Dispropotination Reaction
a. Strong bases cause disproportination into Co(+2) and Co(-1)
2Co(CO)8 + 12NH3  2[Co(NH3)6][Co(CO)4]2 + 8CO
b. With isocyanides it gives penta-co-ordinate cobalt(I) cation.
Co2(CO)8 + 5CNR  [Co(CNR)5][Co(CO)4]+4CO
(ii) Dodecarbonyltetra Cobalt, [CO4(CO)12]
Preparation
1. It is prepared by heating Co2(CO)8 at 60oC.
2. It may also be obtained by oxidizing cobalt carbonyl hydride
below -26oC.
Properties:
i. It is black crystalline solid.
ii. It is very unstable easily oxidized by air and can be
recrystallized from hot benzene.
149
C. Carbonyls Of Nickel.
(i) Nickel Tetracarbonyl, Ni(CO)4
Preparation
i. Ni(CO)4 can be prepared by the action of CO on reduced nickel
at 30-50oC.
Ni + 4CO  Ni(CO)4
ii. When Nickel iodide is heated with CO in the presence of a
halogen acceptor, nickel carbonyl is formed.
NiI2 + 4CO  Ni(CO)4 + I2
Properties
i. It is colourless liquid, m.p. = -23oC, b.p. = 43oC
ii. It has no solubility in water but dissolved in organic solvents.
iii. It decomposes at 180o-200oC in to nickel and CO.
C
Ni(CO)4 180

 Ni + 4CO
3
iv. It reacts with H2SO4 and form NiSO4
Ni(CO)4 + H2SO4  NiSO4 + H2 + 4CO
v. It reacts with Ba(OH)2 and gives BaCO3
Ni(CO)4 + Ba(OH)2  H2Ni(CO)3 + BaCO3
D. Carbonyls of Ruthenium
It forms three carbonyls :
a. Ruthenium Pentacarbonyl, Ru(CO)5
Preparation
i. It is prepared by the action of CO and reduced ruthenium at
200oC and 200 atm. pressures
Ru + 5CO  Ru(CO)5
150
Properties
i. It is colourless soluble liquid having m.p. = -22oC
ii. It has no solubility in water but is soluble in alcohol,
benzene and CHCl3.
iii. It
undergoes
decomposition
to
give
Ru2(CO)9
and
Ru3(CO)12.
iv. It reacts with halogen to yield Ru(CO)Br and CO.
v. It is photosensitive and yields ruthenium enneacarbonyl.
b. Ruthenium Enneacarbonyl, Ru2(CO)9
It is prepared by exposing pentacarbonyl to u.v. radiation. It
forms yellow monoclinic crystals. It is volatile; it is less stable
towards heat, with iodine. It yield, Ru (CO)2I2.
c. Ru3(CO)12
It is prepared in small quantities along with Ru2(CO)9 when
Ru(CO)5 is heated at 50oC or by exposing Ru(CO)5 to u.v. light.
It is a green crystalline solid.
E.
Carbonyl Of Osmium
It forms two carbonyls:
a. Osmium pentacarbonyl. Os2(CO)5
It is a colourless having m.p.-15oC. It is obtained.
i. by the action of CO on OsI3 at 120oC and 200 atm. pressure
in the presence of copper.
ii. by the action of CO on OsO4 at 100oC and 50 atm. pressure.
OsO4 + 9CO  Os(CO)5 + 4CO2
b. Osmium eneacarbonyl, Os2(CO)9
It is a yellow crystalline solid. It is prepared by the reaction
of OsI3 with CO in the presence of copper, it is more stable
towards heat than Ru4(CO)9. It melts at 224oC and sublimes
without decomposition.
151
F.
Carbonyl Of Iridium
It forms 2 carbonyls:
a. Iridium Octacarbonyl, Ir2(CO)8
It is prepared by the reaction of either KIr2Br6 or KIr2Br6 or
KIr2I6 with CO at 200oC and 200 atm. pressures.
It is yellow crystalline solid having m.p. 160oC.
b. Iridium Dodecarbonyl, Ir4(CO)12
It forms orange yellow rhombohedra crystals which
decomposes at 200oC. It is prepared by treating Irl3 with CO
under pressure.
G. Carbonyl of Platinum
Preparation
i.
When CO is passed over PtCl2 at 250oC, PtCl2(CO) and
2PtCl23CO are obtained on heating these yield PtCl2(CO)2.
3PtCl2 + 5CO  PtCl2.2CO + 2PtCl2.3C0
Properties
These carbonyls are decomposed by water and HCl.
PtCl2.CO + H2O  Pt + 2HCl + CO2
PtCl2.CO + H2O  Pt + 2HCl + CO2 + CO
PtCl2.CO + HCl  H[PtCl3.CO]
PtCl2.CO + HCl  H[PtCl3.CO] + CO
5.2.4 Structure of Metal carbonyls
1. Effective Atomic Number Rule:
The structure of CO is :C:O:
It is probable that the lone pair of electrons on the carbon
atom can be used by forming a dative bond with certain metals
152
(MC  O, Thus (MC  O) types of bonds were assumed to be
present in metal carbonyls. In the formation of MC  O bonds,
the electrons are supplied by the molecules of CO and the metal
atom is thus said to have zero-valency. The Number of molecules
of carbon mono-oxide which can unite with one atom of the metal
is controlled by the tendency of the metal atom to acquire the
E.A.N. of the next inert gas. For the stable nonnumeric carbonyl.
E.A.N. = m + 2y = G
Where
M = Atomic number of the metal M
Y = No. of CO molecules
G = At. No. of next inert gas
Carbonyls
Cr(CO)6
Atomic
Number of
the metal
24
Number of electron E.A.N. Succeeding
contributed by CO
inert gas
groups
12
36
Kr(36)
Fe(CO)5
26
10
36
Kr(36)
Ni(CO)4
28
8
36
Kr(36)
Mo(CO)6
42
12
54
Xe(54)
Ru(CO)5
44
10
54
Xe(54)
W(CO)6
74
12
86
Rn(86)
Os(CO)5
76
10
86
Rn(86)
on the basis of E.A.N. rule it can be explained why Ni atom fails to
form a hexacarbonyl Ni(CO))6 because EAN or Ni atom in Ni(Co)6
would be equal to 28 + 2 x 6 = 40. Which is not the atomic number
of any of the noble gases. Mononuclear carbonyls having the
metallic atom with odd At. No. V(CO)6 and Mn(CO)5 & Co(CO)4
are the example of such carbonyls. They do not obey EAN rule
V = 23e6CO = 12 eV(CO)6= 35e-
Mn = 25 e5CO = 10 eMn(CO)5= 35e153
Co = 27 e4CO= 8 eCo(CO)4= 35e-
Therefore the metals with odd atomic number cannot form
monocular carbonyls but forms polynuclear carbonyls for example,
Mn(25) and Co(27) form polynuclear carbonyls.
2. Polynuclear Carbonyls:
Sidgwick and Bailey gave the general formula for
polynuclear carbonyls.
G
Where
X m  2y
 X 1
X
G = The At. No. of next Inert Gas.
M = The At. No. of metal atom.
Y = The No. of CO molecules in one
molecule of the carbonyl.
Mn2(CO)10, CO2(CO)8 etc. obey the E.A.N. rule, their E.A.N. per
atom of metal is 36.
For example: E.A.N. of Mn2(CO)10 may be calculated as:
Electrons from 2Mn Atom
=
25 x 2 = 50
Electron from 10CO molecules =
10 x 2 = 20
Electrons from one Mn-Mn Bond=
1x2=2
Total =
72
E.A.N. for one Mn atom = 72/2 = 36
The formation of binuclear carbonyls having metal atoms
with odd atomic number can also be explained on the basis of 18electron rule as shown below for Co2(CO)8.
Co2(CO)8
2Co = 2 x 9e-
=
8Co = 2 x 8e-
18e=
Co-Co bond = 1x 2e- = 2eCo2(CO)8 = 36e
Electrons on one Co atom = 18e154
16e-
Drawback
X-Ray diffraction method shows that the bonds are
intermediate between the M-C = 0 and M=C=0 states, i.e. there is
some double bond character in M-CO. The EAN rule does not explain
double bond character. This is explained by both MOT and VBT.
3. Molecular Orbital Approach
According to the M.O.T. carbon and Oxygen atom undergo
overlapping to form bonds in CO as follows
i.
2 sp hybrid orbital of carbon and 2px of oxygen overlap to
form a localised   bond.
ii.
2py of carbon and 2py of oxygen overlap to form a π-bond.
iii. 2pz of carbon and 2 pz of Oxygen overlap to form another πbond.
iv. There will be 2 non-bonding electrons in the 2sp hybrid
orbital of carbon.
v.
There will be 2 non-bonding electrons in 2s atomic orbital of
oxygen.
vi. There will be no electron in the anit-bonding molecular
orbitals, formed as result of  anti π overlapping.
As the total No. of bonding electrons is six and that of
antibonding electrons nil, bond order of the molecule is three.
Hence, the No. of bonds between carbon and oxygen atoms in CO
molecules is 3, one  and two π.
The lone pair of electrons on carbon could be expected to
form a strong dative bond (  ) due to the electron density
remaining close to the nucleus of the carbon atom. As metal atomcarbon mono-oxide bonds are readily formed in metal carbonyls. It
155
is expected that there is some additional bonding mechanism in the
formation of metal-carbon monooxide bonds in the metal
carbonyls.
Mechanism
1.
Firstly, there is a dative overlapping of filled carbon  -orbital
i.e. 2sp hybrid orbital with an empty metal  -orbital (MCO)
as in the figure 5.1.
-
m
+
+ C = 0 :
-
M
+
C = O:
Fig. 5.1 L - M  bonding
2.
Secondly, there is a dative overlapping of a filled d-orbitals of
metal with empty antibonding p-orbital of the carbon atom
(MCO), resulting in the formation of a dative π bond. The
shaded portion in figure, indicate the filled orbitals, whereas
empty portions indicate vacant orbitals. i.e. having no
electrons.
As there is a drift of metal electrons into CO (MCO)
orbitals will tend to make the CO as a whole negative and at
the same time there is a drift of electrons from CO to the metal
(MCO) to make CO positive. Thus enhancing the acceptor
strength of the π bond formation and vice versa.
Fig. 5.2(a) dπ - Pπ back bonding
156
Fig. 5.2(b) M - CO  and π bonding
3. Valence Bond Method: Monocular Carbonyls
In this method, the molecule may be represented by
resonance structures.
M -+ C – O  M = C = O
with a large amount of the double bond character, it is this structure
that account for their stability. From either the molecular orbital or
the valency bond view point, back donation is seen in both.
Structure of Ni(C)4
1.
The vapour density of nickel carbonyl and the freezing points
of its solution in benzene indicate the molecular formula to be
Ni(CO)4.
2.
Electron diffraction studies shows that Ni (CO)4 molecule has
tetrahedral shape with Ni-C-O linear units. Figure shows that
the Ni-C bond length in this molecule is 1.50Ao which is
shorter by 0.32Ao in comparison in Ni-C single bond length
(=1.82Ao) found in carbonyls. The C-O bond length in this
carbonyl has been found to equal to 1.15 Ao. Which is larger
that the C-O bond length in CO molecule (=1.128Ao) (Fig.
5.3)
Fig. 5.3 L Tetrahedral structure of Ni (CO)4 molecule
157
3.
Raman Spectra shows that nickel atom in the nicle carbonyl
must be tetrahedrally hybridised as in the figure 5.3.
Titrahedral shape of Ni (CO)4 arises due to Sp3 hybridisation of
Ni-atom. Which is diamagnetic, all the ten electrons present in the
valence shell of Ni atom are paired in 3d orbitals. Thus the valence
shell configuration of Ni atom in Ni (CO)4 molecule becomes 3d10 4SO
CO  Ni bond is caused by the overlap between the empty sp3 hybrid
orbital on Ni-atom and doubly filled sp hybrid orbital on C atom in
CO molecule, as in the figure 5.3 (b)
Because of the formation of 4 OC  M bonds, a large negative
charge gets accumulated on central Ni atom. Pauling suggested that
the double bonding occurs with the back donation of d-electron from
Ni atom to CO ligands to such an extent that electroneutrality
principle is obeyed. According to which the electron pair is not shared
equally between Ni and C-atoms of CO ligand but gets attracted more
strongly by C-atom which prevents the accumulation of negative
charge on Ni-atom, in keeping with the greater electronegativity of Catom compared to Ni atom.
Evidences:
1.
The above structure (Fig. 5.3) is supported by the following
reactions:
i.
When an alcoholic solution of the carbonyl is treated with
orthophenanthroline to yield a stable ruby-red compound
Ni(CO)2 phen. It confirms that two C=O groups of Ni(CO) 4
are replaced by one molecule of phenanthroline.
ii.
Similarly the reactions of Ni(CO)4 with diarsine indicates the
two C=O groups are replaced and remaining two are retained.
158
2.
i.
Structure of Fe(CO)5: The various evidences are:
The vapour density and the freezing points of benzene
solution shows that its molecular formula is Fe(CO)5.
ii.
Electron diffraction, Raman and I.R. spectra shows that it has
trigonal bipyramidal shape and Fe-C axial bond and Fe-C
basal bond lengths are equal to 1.797Ao and 1.842Ao
respectively. It has dsp3 hybridisation of Fe atom (Fig. 5.4(c).
III. Molecule is Diamagnetic and the distance Fe-C is 1.84Ao(Fig.
5.4).
Fig. 5.4 (a) : Structure of Fe(CO)5
Structure of Cr(CO)6
It has octahedral configuration. The internuclear bond
lengths are:
Cr-C
Cr-O
C-O
1.92
3.08
1.16Ao
According to old concept when chromium forms Cr(CO)6
one electron of 4s orbital missing and three 4p orbitals become
empty which are hybridised to form six d 2sp3-hybrid orbitals six
molecules of CO donate a lone pair of electrons each to six vacant
hybrid orbitals to form six CrCO  -bonds as shown in the Figure
5.5. Therefore Cr(CO)6 molecule is diamagnetic in nature and
octahedral in geometry.
159
When chromium atoms form chromium carbonyl [Cr(CO)6]
the metal atom exhibits d2sp3 hybridisation. Out of 6 d2sp3 hybrids
three hybrid orbitals are half filled and three hybrid orbitals are
empty. Three electrons remain in 3d orbitals as shown in the figure.
5.5(a) and (b). The Bond Structure of Cr(CO)6 Shows 2 kinds of
bonds between Cr and Co.

(a) Simple Covalent Bonds Cr-C  0 (b) Double bonds Cr 
c=O
In the resultant resonance structure all Cr-C bonds have been
identical, each of the 6 CO groups get linked to the metal atom by a
 bond are constructed from the d-orbitals of the metal atom. CO
(groups I) are bound to the metal atoms by simple ionic bonds.
Fig. 5.5 Structure of Cr(CO)6
Hence these CO groups are replaceable by any other
molecule capable of donating lone pair of electrons to the metal
atom where as CO (groups II) are not replaceable.
In the same way structure of MO(CO)6 and W(CO)6 can be
explained. Various internuclear bond lengths of these carbonyls are
as under:
TABLE : INTERNUCLEAR BOND LENGTHS
Metal
M-C(A)
M-O(A)
C-O(A)
Cr
1.916
3.98
1.171
Mo
2.063
3.23
1.145
W
2.06
3.19
1.148
160
Thus, the bond structure of Cr(CO)6 shows 2 kinds of bonds
between Cr and CO.
i.
Simple covalent bonds
ii.
Double bond
Cr-C  0
(I)
Cr = 0
(II)
Structure of Polynuclear Carbonyls
These crabonyls obeys EAN rule, if two electrons from each
metal metal bond present in these carbonysls are included in
calculating the electrons per metal atom; eg metal-metal bonding is
evident in Mn2(CO)10 as in Figure. 5.6
Structure of Dinuclear Carbonyls
(a) (i) Mn2(CO)10 : Its structrue is shown in Fig 5.6.
161
Fig. 5.6
(ii) Structure of Fe2(CO)8
I.R. and X-Ray study show that in this molecule each Fe
atom is directly linked with the other Fe atom by a S-bond (Fe-Fe Sbond) to three bridging carbonyl gropus (>C = 0) by a  bond (Fe-C
 bond) and to three terminal carbonyl gropus (-C  = 0) by a co-
ordinate bond (FeC co-ordinate bond). The presence of Fe-Fe bond
is supported by the diamagnetic character of Fe2(CO)9 molecule. FeFe bond distance has been found to be equal to 2.46A o. The terminal
C-O bond distances from the structure given in figure. (5.7)
162
The co-ordination number of each Fe atom is not equal to 6
but equal to 7.
Fig. 5.7: Structure of Fe2(CO)8
Similarly structures of Co2(CO)8 can be represented as in Fig. 5.8
Fig. 5.8: Structure of Fe2(CO)8
(b) Structure of Trinuclear Carbonyls
Os3(CO)12 and Ru3(CO)12 possess similar structure (Fig.
5.9a) where as Fe3(CO)12 has a different structure 5.9(b). Os and Ru
molecules do not have any bridging CO group (Fig. 5.9 (a)). In
163
Fe3(CO)12 each of the two Fe atoms is linked with three terminal CO
groups, two bridging CO groups and third Fe atom is linked with four
terminal CO groups and to each of two Fe atoms.(Fig. 5.9 (b). It is
also shown by a structure similar to Fig. 5.9.
Structure of Fe3(CO)12
According to old concept each iron atoms gets hybridized
trigonal bipyramidally (dsp3). The three trigonal bipyramides get
arranged in such a manner so that the carbonyl groups at two of the
equatorial apices of each bipyramid and held in common by two
bipyramides dxz and dyz orbitals are available to form Fe-Fe bonds. It
is solid. The three Fe atom get situated at the corner of an isosceles
triangle and the twelve CO arranged at the twelve CO arranged at the
vertices of an icosahedra. Two Fe-Fe bond lengths are 2.698 Ao and
one Fe-Fe bond length is 2.56Ao. (Fig. 5.9)
Fig. 5.9 Fe3(CO)12 Complex
Fig. 5.9 (a) M3 (CO)12 Structure
164
Fig. 5.9 (b) Structure of Fe3 (CO)12
(c) Tetra and Hexanuclear Carbonyls
On the similar grounds structures of tetranuclear carbonyls, such as
M4(CO)12 [M = Co, Rh, Ir] and hexanuclear carbonyls, M6(CO)16 eg.
Rh6(CO)16 can be represented as in Figs. 5.10 and 5.11 respectively.
Some heteronuclear carbonyls are also known e.g. Mn2Fe(CO)14 is
shown in Fig. 5.12.
(a)
(b)
Fig. 5.10 (a) Structure of Ir4(CO)12
(b) M4(CO)12; M = Co or Rh
165
Fig. 5.11 Structure of Rh6(CO)16
Fig. 5.12 Structure of Mn2Fe(CO)14
5.2.5 Vibrational Spectra
IR spectra give important information regarding nature of carbonyl
groups present in metal carbonyl complexes. We can differentiate
between the terminal carbonyl e.g. in Mn2(CO)10 and bridging carbonyl
groups, as in Co2(CO)8.
Metal-carbon distances in Fe2(CO)9 and Co2(CO)8 fall into two
groups, metal-bridging carbonyl distances being about 0-1 A longer than
metal-terminal carbonyl distances. Such a difference is compatible with
the concept of two-electron donation by terminal carbonyls, and oneelectron donation (to each of two metal atoms) by bridging carbonyls,
through the possible existence of  bonds of different strengths makes
166
quantitative interpretation impossible. That the extent of  bonding to
terminal and bridging carbonyls is different is clearly shown by carbonyl
stretching frequencies. Carbon monoxide itself has stretching frequency
of 2143 cm-1; neutral metal carbonyls known to have no bridging
carbonyl groups have stretching frequencies in the range 2125-2000 cm-1;
and Fe2(CO)9 and Co2(CO)8, in addition to showing bands in this region,
also show carbonyl absorption at 1830 and 1860 cm-1 respectively. In
general, carbonyl absorption in the 1900-1800 cm-1 region is indicative of
the presence of bridging carbonyl groups in uncharged species, though
the presence of other groups may result in the lowering of the stretching
frequencies of terminal carbonyl groups into this region (in carbonylate
anions such as [Co(CO)4]- and [Fe(CO)4]2- very low carbonyl stretching
frequencies of 1883 and 1788 cm-1 respectively result from the strong
metal-carbon   bonding which stabilizes the low oxidation state of the
metal. In a few neutral species believed to contain carbonyl groups
bonded to three metal atoms, stretching frequencies of 1800 cm-1 or less
are found.
Thus, in summary the terminal carbonyl absorption is obtained in
the range of 2125-2000 cm-1, while bridging carbonyl frequency is
obtained in the 1900-1800 cm-1 region. While, the strong metal carbon
 bonding is indicated by very low carbonyl structing frequencies of 1883
and 1788 cm-1 respectively.
167
Check Your Progress-1
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
(a) Generally metal carbonls involve...............................  -and .........
.....................  -bonding between CO and metal atom.
(b)
Metal
atoms
with
even
number
of
electrons
easily
form............carbonyl, but the metal ions with odd number of
electrons give.......................or...............................
(c) I.R. spectra of carbonyls show C-O stretching frequency
at..........................cm-1 for the terminal carbonyl group and
at................cm-1 for the bridging carbonyl group. The M-CO π
bond is indicated by the absorption at...................and.............Cm-1
respectively.
(d) While Mn2(CO)10 has........................bridging carbonyl group,
Fe2(CO)9 has................................................................................
5.3
METAL NITROSYLS
Nitrosyls are the compounds in which the nitrogen of the nitrosyl
group is directly bonded to the atoms or ions, or the compounds
containing nitric oxide group are called nitrosyl compound. NO molecule
is an odd electron molecule having an unpaired electron, it readily unites
with other elements by direct addition to form nitrosyl compounds. Nitric
oxide form nitrosyl compound by the following 3 ways:
i.
A positive ion, NO+ is formed due to the loss of an electron
which then combine with atom or molecule (:N:::O)+ or (:N  O:)+
168
ii.
A negative ion NO- is formed due to the gain of an electron from
some electropositive metal and it has structure as below:
(:N:: O)- or (:N=O:)-
iii. NO may act as a co-ordinating group through the donation of an
electron pair, such behaviour involve neutral molecule or NO + or
NO- group.
The electronic configuration of NO group is so flexible that it is
rather impossible to write its any one configuration in metal nitrosyls.
However in all nitrosyls nitrogen atom is linked with the metal possible
modes are given below:
I


:N
:N
:O:
Links
:O:
Links
II
III

with


bonds

(neutral)
bonds
D
:N

V
D
:N

:O:
with Links
and
IV

:O:
with Links
:N
:O:with Accepts
dative bonds
dative bonds electron
(cationic)
(cationic)
(cationic)
from metal
(anionic)
Mode (I) is rarely seen, while (II), (III) and (IV) modes are formed
after transferring one electron to the metal atom. Out of these modes (II)
and (III) are similar to the carbonyl group linked with a metal atom.
Mode (V) is seen only in a few complexes only, e.g. [Co(CN)5NO].
5.3.1 Neutral NO and NO- Complexes
As has been pointed out metal complexes of neutral NO and its
anion, NO- are very rare. Fe(NO)2(CO)2 is supposed to be the important
example of metal complex with neutral nitric oxide molecule. This is
prepared by the action of nitric oxide on Fe(CO)5. During the reaction,
169
NO replaces neutral CO, hence it is supposed to be a complex of neutral
NO. However, the experimental evidences are not supportive.
The important examples of anionic NO- are the metal complexes,
formed by the action of nitric oxide with ammonical solution of Cobalt
(II) salts, with the general formula [Co)NH3)5NO]X2. Two series of
isomeric complexes are formed one having black colour, while the other
one has red colour.
The black series contains the monomeric cation [Co(NH3)3NO]2+,
in which a very low N-O stretching frequency of 1170 cm-1 and a long NO bond (variously reported as 1.26 or 1.41A) suggests the presence of
NO-. The red series are derivatives of hyponitrite, the structure of the
diametric cation being.
Similarly [Co(CN)5NO]3- anion is also supposed to be a complex of
NO- anion, since it gives NO- stretching frequency at 1150 cm-1.
5.3.2 Complexes of NO+
Most complexes of nitric oxide and transition metals are best
considered to be those of the NO+ ion, three electrons being transferred to
the metal atom: M-N back π-bonding then takes place in exactly the same
way as for carbon monoxide. Because of its positive charge, however,
coordinated NO is a better π-acceptor than coordinated CO, and the N-O
170
stretching frequency in complexes of NO+ is some 300-500 cm-1 lower
than that in salts such as NO+BF4-.
Two NO+ derivatives of iron may be mentioned briefly here. The
species formed in the brown-ring test for nitrate is [Fe(H2O)5NO]2+. The
equilibrium
[Fe(H2O)6]2+ + NO  [Fe(H2O)5NO]2+ + H2O
is reversible, and the brown complex may be destroyed by blowing
nitrogen through the solution to remove nitric oxide. In this species the
N-O stretching frequency is 1745 cm-1, and the magnetic moment is 3.9
B.M., corresponding to the presence of three unpaired electrons;
formally, therefore, the ion is a high-spin d7 complex of Fe1 and NO+, but
the N-O stretching frequency indicates very strong π-bonding and the
intense brown colour strongly suggests Fe1-NO+ charge transfer.
5.3.3 Pure Nitrosyl Complexes
Pure nitrosyl complexes of M(NO)4 formula have been reported.
Important complexes in this series are Fe(NO)4, Ru(NO)4 and Co(NO)4.
In addition to this trinitrosyl cobalt, Co(NO)3 has also been reported.
Fe(NO)4 is prepared by the action of nitric oxide under pressure
and below 45oC temperature on Fe(CO)5. While M(NO)4 nitrosyls of Ru
and Co are prepared by the same method using Ru 2(CO)9 and Co2(CO)8
respectively.
Fe(NO)4 is a black crystalline substance which decomposes in to
Fe(NO) and Fe(NO)2. The structure of tetranirtrosyl iron, Fe(NO)4 has
been shown tetrahedral, while that of trinitrosyl cobalt, Co(NO) 3
pyramidal. Nitric oxide links with iron, following II mode, as a three
electron donor and results in a strong ML back π-bonding (Fig. 5.13).
171
5.3.4 Nitrosyl Carbonyl Complexes
Mononuclear nitrosyl carbonyls are restricted to the following
compounds; Co(NO)(CO)3, Fe(NO)2(CO)2, Mn(NO)3CO and Co(NO)3
(isoelectronic with Ni(CO)4; Mn(NO)(CO)4 (isoelectronic with Fe(CO)3;
and V(NO)(CO)5 (isoelectronic with Cr(CO)6). In addition a binuclear
species Mn2(NO)2(CO)7 (isoelectronic with Fe2(CO)9 and a number of
nitrosyl complexes containing organic groups or triphenylphophine as
substituents have been prepared. Nitric oxide displaces carbon monoxide
from V(CO)6, (Ph3P)2Mn2(CO)8, Fe2(CO)9 and Co2(CO)8 to give
V(NO)(CO)5,
Mn(NO)(CO)4,
Fe(NO)2(CO)2
and
Co(No)(CO)3
respectively; the further action of nitric oxide on the manganese and
cobalt compounds yields Mn(NO)3(CO) and Co(NO)3. All of these
substances are solids of low melting point or liquids which are thermally
rather unstable and are decomposed by air and by water. In the reaction of
Fe(NO)2(CO)2 with alkali in methanol, [Fe(NO)(CO)3]- is formed, but
under comparable conditions Co(NO)(CO)3 gives [Co)CO)4]-, Co(OH)2
and other cobalt-free products.
The limited evidence available is consistent with tetrahedral
structures for Fe(NO)2(CO)2 (Fig. 5.14) and Co(NO)(CO)3 and a trigonal
bipyamidal structure (with NO in the equatorial plane) for Mn(NO)(CO)4
(Fig. 5.15); (Ph3P)2Mn(NO)(CO)2 also has a trigonal bipyramidal
structure, the two triphenylphosphine molecules occupying the apical
positions. Since Co(NO)3 shows two N-O stretching frequencies in the
infrared, it must be pyramidal rather than planar, but the detailed structure
is not known.
5.3.5 Nirtosyl Halide Complexes
Volatile diamagnetic nitrosyl halides of formula Fe(NO) 3X are
formed by the action of nitric oxide on iron carbonyl halides in the
172
presence of finely divided iron as a halogen-acceptor. These readily lose
NO to give [Fe(NO)2X]2, in which the halogen atoms act as bridges.
Analogous compounds of cobalt and nickel may be formed by reactions
similar to those involved in the high pressure synthesis of carbonyls; for
example,
CoX2 + Co + 4NO  2Co(NO)2X
4NiI2 + 2Zn + 8NO  2[Ni(NO)I]4 + 2ZnI2
The ease of formation of these compounds increases in the
sequences Ni < Co < Fe and X = Cl < Br < I. Nitrosyl chloride and nickel
carbonyl in liquid hydrogen chloride, on the other hand, give Ni(NO) 2Cl2,
which is probably monomeric and tetrahedral. Nitrosyl halides are also
formed by some metals which, so far as is known, do not form nitrosyls
or nitrosyl carbonyls. Thus molybdenum and tungsten (but not chromium)
carbonyls react with nitrosyl chloride:
M(CO)6 + 2NOCl
20
M(NO)2Cl2 + 6CO
CH 2 Cl 2
Palladium (II) chloride in methanolic solution yields Pd(NO) 2Cl2,
and nitrosyl halide molecules or anions are formed also by several other
transition metals.
5.3.6 Nirtoso Cyanide Complexes
Sodium nitropursside is also a complex resulted from the
coordination of NO+.
Sodium nitroprusside [nitrosopentacyano-ferrate (II)] is prepared
by the action of nitric acid or sodium nitrite on the hexacyanoferrate (II).
In the former process the overall reactions is
[Fe(CN)6]4- + 4H+ + NO3-  [Fe(CN)5NO]2- + CO2 + NH4+
173
In the latter process, two successive equilibria are involved:
[Fe(CN)6]4- + NO2-  [Fe(CN)5NO2]4- + CN[Fe(CN)5NO2]4- + H2O  [Fe(CN)5NO]2- + 2OHThese are driven to completion by adding barium chloride to the
reaction mixture and blowing a current of carbon dioxide through the hot
solution to remove the hydrogen cyanide liberated by the reaction
2[Fe(CN)6]4-+2NO2-+3Ba2++3CO2+ H2O  2[Fe(CN)5NO]2-+
2HCN + 3BaCO3
The formulation of the complex anion as a NO+ derivative of iron
(II) is supported by its diamagnetism, a N-O stretching frequency of 1939
Cm-1 and a N-O distance of 1.13A. The purple colour obtained from
nitrosopentacyanoferrate
(II)
and
sulphide
is
due
to
the
ion
[Fe(CN)5(NOS)]4- analogous to [Fe(CN)5NO2]4-.
Structure of Nitrosyl Co-ordination Compounds:
If we compare the electronic structure of NO with CO, it is
observed that NO has an additional electron in antibonding π M.O.,
which may be readily lost to form the nitrosonium ion, NO+.
The additional electron present in π molecular orbital of NO can be
supplied to metal atom thus increasing its effective number by one unit
and neutral No is itself converted into NO+ ion. Then, this NO+ is coordinated through nitrogen with the metal atom by donating its lone pair
to the metal.
1.
Cobalt atom may increase its E.A.N. from 27 to 28 by accepting an
additional electron from a neutral molecule of NO:
Co + NO  Co- + NO+ cobalt ion may then combine with
one NO+ group and 3 CO molecules to form stable compounds Co 174
+ NO+ + 3CO  Co(NO)(CO)3 In this compound the E.A.N. of Co
is, 27+1+2+6 = 36 of stable Krypton.
2.
Similarly the formation of Fe(NO)2(CO)2 can be explained.
Sidgwick gave the electronic structure of metallic nitrosyls as
belowM+  (:N: ::O:+) or M2- - N+  O+
The accumulation of charge on the central atom favours
strong π-bond formation with the attached groups.
Thus the most of metal nitrosyls are formed by donation
from the (NO)+ to the metal atom with the M-O back bonding in a
manner analogous to M-C bond in carbonyl it is known as three
electron donor
M + NO  M- + NO+ M2- - N+ = 0+
In terms of M.O.T. the hybrid orbital on N atom having a
lone pair [(sp)2N lone pair] overlaps suitable vacant hybrid orbital
on M ion (sp3 in tetrahedral or d2sp3 in octahedral) to form ON+ 
M-  -bond and the empty π2* or π1* M.O. will overlap with the
filled d-orbitals to form M- NO+ π bond. This type of overlap
transfers some charge from M- ion to NO+ ion. The molecule of
NO is a resonance structure of the following forms:
π
π
N 
O:  N – O  -N 
O+:
On this basis resonance structures of NO, the metallic
nitrosyls may be represented as:
π
π
π
π
M--N 
O: M- N 
O M 
N 
O: M--N-O
Nitric oxide is a paramagnetic molecule with an electron in
175
an anti-bonding orbital. This electron is relatively easily lost with
formation of the NO+ ion and an increase in the N-O stretching
frequency from 1878 cm-1 in NO to 2200-2400 cm-1 in nitrosonium
salts.
Structure of various groups of nitrosyl complexes are shown
in Fig. 5.13 to 5.20
Fig. 5.13 Structure of Fe (NO)4
Fig. 5.14 Structure of Fe(CO)2 (NO)2
Fig. 5.15: Structure of [Mn(NO) (CO)4]
176
Fig. 5.16: Structure of [Fe(NO)2I]2
Fig. 5.17: Structure of Fe(NO)3Cl
Fig. 5.18: Structure of [Ni (NO)I]4
177
Fig. 5.19: Red salt of Diethyl Ester of [FeS2(NO)4]2-
Fig. 5.20: Anion of Red salt of [Fe4(NO)7S3]
5.4
DINITROGEN COMPLEXES
In 1965, Allen and Senoff obtained salts containing the
[Ru(NH3)5N2]2+ cation by the action of hydrazine hydrate on various
compounds of tri- and tetrapositive ruthenium, amongst them ruthenium
trichloride and ammonium hexachlororuthenate (IV).
Thus, these substances (often called nitrogenyl or dinitrogen
complexes, to distinguish them from those containing the nitride ion)
have been known for only a few years.
178
Many other complexes containing one or two (but not, so far,
more) molecules of coordinated nitrogen have now been prepared, and it
is clear that N2 acts as a  -donor and π-acceptor in the same way as
isoelectromic CO, though the complexes formed are much less stable
than carbonyls. Much of the interest in this field centres on the possibility
of developing new methods for nitrogen fixation; up to the present time,
however, no method has been found for the reduction of nitrogen in the
complexes described here (though this has been achieved by systems
involving an organ titanium complex under powerfully reducing
conditions).
Most,
though
not
all,
nitrogenyl
complexes
have
triphenylphosphine and halide or hydride as other ligands in the complex.
The following examples illustrate methods for their preparation.
(a) The action of nitrogen on a metal complex: for example,
CoCl2 + Ph3P
NaBH 4
N
(Ph3P)3CoH3 
(Ph3P)3CoH(N2)
EtOH
3
N
[Ru(NH3)5H2O]2+ 
[Ru(NH3)5(N2)]2+
3
(b) Another method of preparation of dinitrogen complex is the
reaction of coordinated azide:
[Ru(NH3)5Cl]2+ N3-
MeSO3 H
[Ru(NH3)5N2]2+
NH 3
(c) Similarly reaction of (Ph3P)2Ir(CO)Cl with RCON3:
(Ph3P)2Ir(CO)Cl + RCON3  Ir(PPh3)2(CO)(Cl)(N2.NCOR)
CHCl3 / EtOH
(Ph3P)2IrCl(N2)
179
(d) Reaction of coordinated NH3 with HNO2:
[Os(NH3)5(N2)]2+ + HNO2  [Os(NH3)4(N2)2]2+ + 2H2O
The most stable dinitrogen complexes are those of heavier
members of iron and cobalt groups. Some are unaffected by dry air and
can be heated to 100-200oC without decomposition. Most are rapidly
oxidised by air and decompose on heating gently.
The orange solid (Ph3P)3CoH(N2+) shows reversible displacement
with hydrogen, ethylene or ammonia. Some of the reactions of
(Ph3P)2IrCl(N2) (yellow solid) are as follows:
(Ph3P)2IrCl(N2) + Ph3P  (Ph3P)3IrCl + N2
(Ph3P)2IrCl(N2) + HCl  (Ph3P)3IrHCl2 + N2
(Ph3P)2IrCl(N2) + CO  (Ph3P)2Ir(CO)Cl + N2
Dinitrogen complexes show an asymmetric IR N  N stretching
frequency in the range 2230-1920 Cm-1(Raman stretching frequency in
N2 is 2331 cm-1).
In metal complexes dinitrogen either has a terminal position or as a
bridge:
N
M–N–N
N
M–N–N–N M
M
N
Terminal
Structures
of
the
M
N
Bridging
two
important
dinitrogen
complexes
[Ru(NH3)5N2Ru(NH3)5]4+ and [Sm(N5C5Me5)2I2(N2)] are shown in Fig.
5.21 and 5.22 respectively.
180
Fig. 5.21: Structure of [Ru(NH3)5(N2)Ru(NH3)5]
Fig. 5.22: Structure of [Sm(N5C5Me5)2I2(N2)]
5.4.1 Fixation of Nirtrogen
Dinatrogen complexes while show possibility of developing new
methods for nitrogen fixation, they also help in the understanding of the
probable mechanism of biological fixation of nitrogen.
An important enzyme-system is related with the atmospheric
fixation of nitrogen; which involves an important step in nitrogen-cycle
and is responsible for supply of nitrogen to the plants growth (e.g. Bluegreen algae, symbiotic bacteria legume)
181
The active enzyme in fixation nitrogen is nitrogenase. In this
enzyme two proteins take part in the reaction. Small protein has
molecular weight of 57000-73000 and contains Fe4S4 group; while the
large protein is a tetramer of molecular weight 220000-240000. It has 2
molybdenum atoms, nearly 30 iron atoms and nearly 30 mobile sulphide
ions. Fe-S group probably functions as redox centre, and the active site
for dinitrogen binding is probably molybdenum atom. (Fig. 5.23)
Fig. 5.23: Fixation of Nitrogen
5.5
DIOXYGEN COMPLEXES
Amongst all the donor atoms oxygen is most important. The donor
ability of oxygen is related with its partial charge; higher is the negative
charge, higher will be the donar ability. Large number of coordination
compounds are available in which oxygen uses one of its two lone pairs
of electrons.
The most important example of dioxygen complexation is
transportation of oxygen in aerobic-organisms through heme and
hemocynin mechanism. Although, hemoglobin and hemocynin are known
since long time for their specific ability of absorption and release of
182
oxygen; but now a number of synthetic compounds have this property,
e.g. Bis (Salicylic) ethylenedimmine cobalt (II).
Heme, protein is the most important group of metallic porphyrin,
which functions as a oxygen carrier in aerobic organism. In the centre of
its porphyrin ring is iron (Fe2+), which is linked with the protein part of
haemoglobin. Heme is very much sensitive for reaction with oxygen and
the reactive oxygen complex, forming an intermediate product, is
converted into Fe(II) porphyrin or Hemin.
As has been shown earlier, heme protein functions as the oxygen
carrier during respiration of aerobic organisms. In this process,
vertebrates use two heme-proteirs: hemoglobin and myoglobin.
Hemoglobin takes dioxygen from lungs or gills and passes it to the
tissues. Where it is stored in myoglobin. The cytochromes present in
tissues, which functions as electron carrier, reacts with dioxygen and
reduces it. The oxidation power of dioxygen is thus used in burring of the
food. In this way during transportation storing and use of dioxygen three
heme proteins play important part; these are hemoglobin, myoglobin and
cytochrome.
5.5.2 Hemoglobin
Hemoglobin is the red pigment of blood. It has two parts: (a) 96%
part of it is a simple, specific protein called globin and (b) 4% remaining
part is the prosthetic group hence:
Globin
Hemoglobin
Heme
Protoporphyrin
Fe(II)
183
It is a globular protein, which is made up of polypeptide chins.
These chains are arranged in a regular tetrahedral form and are linked
with the four rings of pyrole. Molecular weight of hemoglobin is nearly
64500.
Hemoglobin molecule can coordinate with dioxygen without
oxidation of iron. The bonding of iron with dioxygen is so strong that
oxyhemoglobin does not decompose during its transportation in the body.
Still it is so weak that its contact with oxidase decomposes it readily. The
various steps during oxidation of hemoglobin are:
Ist step: Bonding with dioxyen:
IInd step: Bonded dioxygen links with other heme (  -peroxo complex
is formed) :
IIIrd step: Decomposition of  per oxo complex into ferryl complex.
IVth step: Reaction of ferryl Complex with heme to give Hematin:
In living being steps I and IV do not take place, otherwise total
heme would have precipitated as hematin. Apart from other reactions,
184
steps III and IV are checked by sterric hinderance. Thus dioxygen is
carried away by oxyhemoglobin and either stored in oxymyoglobin or
given to cytochromes for use.
In lungs or gills of vertebrates, the following reactions take place:
Hb + 4O2  Hb (O2)4
Hemoglobin Oxyhemoglobin
While in tissues, the reaction that takes place is:
Hb(O2)4 + 4Mb  4Mb(O2) + Hb
Myoglobin
5.6
TERTIARY PHOSPHINE AS LIGAND
Large number of triphenylphosphine and similar substituted metal
carbonyls are known, e.g. Ni(CO).(Ph3P)2. This compound is of great
importance as a catalyst for the polymerisation of olefins and acetylenes
e.g. butadiene to cyclooctadiene and acetylene to benzene and styrene.
Analogous compounds can be obtained by the action of triphenyl
phosphine on iron pentacarbonyl. Similarly dicobalt octacarbonyl gives
two products with Ph3P in 1:1 ratio of Co and Ph3P. One compound is
[Co2(CO)6(PPh3)2] and the other is the salt [Co(CO)3(PPh3)2][Co(CO)4] in
which the cation has the expected trigonal pyramidal structure.
A platinum complex, Pt (CO)2(PPh3)2 can be obtained by the action
of
CO
on
Pt(PPh3)4.
As
a
matter
of
fact,
substitution
of
triphenylphosphine for some of the carbonyl groups greatly enhances the
stability of the compound; thus although Co(CO)4 I is unstable,
(Ph3P)Co(CO)3I can be made by the remarkable reaction.
2CF I
 (Ph3P)Co(CO)3I + I + C2F6
[(Ph3P)Co(CO)3]- 
3
Triphenylphosphine carbonyl halides of rhodium and iridium may
be prepared by interaction of the metal halide (or a complex halide) and
185
triphenylphosphine in a variety of organic solvents, the solvent serving as
the source of the carbonyl group:
Ph P
(NH4)2IrCl6 
 (Ph3P)2Ir(CO)CI
3
Ph P
IrCl3.3H2O 
 (Ph3P)2Ir(CO)CI
3
The product of these reactions- (Vaska's compound) is a highly
reactive complex.
Vaska's compound is a carbonyl halide; and many triphenylphosphine complexes containing rhodium and iridium show similar
reactivity and catalytic activity.
The iridium compounds is remarkable for its reversible uptake of
H2, O2 and SO2 to give crystalline 1:1 adducts which can be decomposed
by lowering the pressure; for example,
O
(Ph3P)2Ir(CO)Cl 
O2Ir(PPh3)2(CO)Cl
2
In the oxygen adduct, oxygen atoms occupy cis octahedral
positions; the O-O distance of 1.30 Ao suggests that Oxygen is present as
O2- rather than O22-. Some of the many other reactions of Vaska's
compounds are shown in Fig. 5.24
Fig. 5.24
186
Check Your Progress-2
Notes :(i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the unit.
A.(i) Most of the metal nitrosyls are formed with..................ion.
However, most of the pure nitrosyl complexes have the general
formula............... (M = ...........................................)
(ii) The various modes of linking of NO are:
(a) ......................................
(b) ......................................
(c) ......................................
(d) ......................................
B.
Fixation of nitrogen involves enzyme..........................., which has
two proteins. The small protein, mol-weight.....................,
contains........................group; while the large protein, mole
weight......................... contains................Mo atoms...................Fe
atoms and............mobile sulphide ions.
C. (i)
Hemoglobin
binds
dioxygen
to
give..........................:
(reaction)................................................................................
(ii)
The product is given to...............................for storage and
(iii) is used by...........................for burning of food.
D. Vaska compound has general formula....................................
187
5.7
LET US SUM UP

Most transition metals form complexes with a wide variety of
unsaturated molecules, such as CO, NO, O2, N2 etc., using ML πbonding, which stabilize these complexes.

CO molecules combine with transition metal atoms (generally in
zero oxidation state) to give series of carbonyls, varying from
mononuclear, di-nuclear, tri-nuclear, tetra-nuclear to hexa-nuclear
carbonyls. In which EAN rule is strictly followed.

Metals with even number of electrons give stable mononuclear
carbonyls; but the metals possessing odd number of electrons do
not form stable mononuclear carbonyls. The shortage of one
electron is compensated by linking with H or Cl or by dimmer
formation, e.g. V(CO)6 forms H[V(CO)6, Na[V(CO)6], [V(CO)6]Cl
or V2(CO)12.

IR spectra give important information regarding the nature of CO
group in the complex. Thus the terminal CO group indicate by the
stretching frequency at 2125-1850cm-1 (or 2125-2000cm-1). While
the bridging CO-group is indicated by the stretching frequency at
1900-1800 cm-1 region. Frequency at 1883 and 1788cm-1
respectively are indicative of strong π-bonding (ML).

Ni(CO)4 is tetrahedral, Fe(CO)5 is TBP, Cr(CO)6 is octahedral
while the di-nuclear, trinuclear, tetranuclear and hexanuclear
carbonyls have structures derived from linking of octahedral in
respective numbers of sharing corners or side or a face.

Nitric oxides combine with transition metals to form coordination
compounds. The general modes of linking may be188


(a)

(b)
(c)
D
:N
:N
:N
:O:
Links with 
bonds
(neutral)
:O:
:O:
Links with  Links with
and  bonds dative bonds
(cationic)
(cationic)



(d)
D
:N

:O:
Links with
dative bonds
(cationic)
Most of the nitrosyl complexes are derived from linking of NO +
(nitrosonium ion).

Pure nitrosyls have general formula M(NO)4 with M = Fe, CO, Ru.
However Co(NO)3 has also been reported.

NO displaces CO from V(CO)6, (Ph3P)2Mn(CO)8, Fe2(CO)9 and
Co2(CO)8 to give V(NO)(CO)5, Mn(NO)(CO)4, Fe(NO)2(CO)2 and
Co(NO)(CO)3 nitrosylcarbonyl complexes respectively. In addition
to these, binuclear species such as Mn2(NO)2(CO)7 are also
formed.

Many dinitrogen complexes have been reported e.g. [Ru(NH3)5N2],
[(NH3)5RuN2Ru(NH3)5], [(Ph3P)2IrCl(N2)], [Os(NH3)4(N2)2] etc.

Dinitrogen complexes while show possibility of developing new
methods for nitrogen fixation, they also help in the understanding
of the probable mechanism of biological fixation of nitrogen.

The active enzyme in fixation of nitrogen is nitrogenase. The
enzyme has two proteins one small (mol wt. 57000-73000) protein
contains Fe4S4 groups; while the large protein (mol wt. 220000240000) is a  2  2 tetramer, which has 2 Mo atoms, nearly 30 Fe
atoms and nearly 30 mobile sulphide ions. Fe-S group functions as
189
a redox centre and the active site for dinitrogen binding is
molybdenum atom.

Amongst all the donar atoms oxygen is most important. This ability
is related with its partial charge.

The most important example of dioxygen complexation is
transportation of oxygen in aerobic organisms, through heme and
hemocynin mechanism.

During respiration of aerobic organisms two heme proteins,
hemoglobin and myoglobin, are used. Hemoglobin takes dioxygen
from lungs or gills and passes it to the tissue where it is stored in
myoglobin. The cytochromes present in tissues use the oxidation
power of dioxygen in burring of food.

Hb + 4O2  Hb(O2)4
Hemoglobin Oxyhemoglobin
Hb(O2)4 + 4Mb  4Mb(O2) + Hb
Myglobin Oxyhemoglobin

Large numbers of triphenyl phosphine and similar substituted
metal carbonyls are known e.g. Ni(CO)(Ph3P)2. This compound is
of great important as a catalyst for polymerisation of olefins and
acetylenes.

Substitution of triphenyl phosphine for some of the carbonyl
groups greatly enhances the stability of the compound.

Most
widely
studied
compound
is
'Vaska
compound'
(Ph3P)2Ir(CO)Cl, which is used for the preparation of large number
of triphenyl phosphine containing complexes.
190
5.8
CHECK YOUR PROGRESS: THE KEY
1.(a) Involve CO  M  and M  CO  -bonding.
(b) Form
mononuclear
carbonyl....................give
dimers
or
mononuclear carbonyls linked with H or Cl.
At 2125-2000 cm-1
(c)
and at 1900-1800 cm-1
at 1883 and 1788 cm-1 respectively
(d) Has no bridging group, Fe2(CO)9 has three bridging groups.
2(A) (i) With NO+ ion
formula M(NO)4 (M = Fe, CO and Ru).
(ii)

(a)

(b)
:N
:N
:O:
Links with 
bonds
(neutral)
:O:
Links with 
and  bonds
(cationic)
(c)


D
:N

(d)
D
:N

:O:
:O:
Links with
Links with
dative bonds dative bonds
(cationic)
(cationic)
B. Enzyme nitrogenase, small protein mol. wt 57000-73000
contains Fe4S4 group
Large protein mol. wt. 220000-240000 contains 2 Mo atoms
30 Fe atoms and 30 mobile sulphide ions.
C.(i) To give Oxohemoglobin
Hb + 4O2  Hb(O2)4
(ii) Myoglobin for stroage
and (iii) by cytochromes
D. Vaska compound has general formula:
(Ph3P)2Ir(CO)Cl
191
Unit - 6
REACTION MECHANISM OF TRANSITION METAL
COMPLEXES-I
Structure
6.0
Introduction.
6.1
Objectives.
6.2
Energy Profile of a Reaction.
6.3
6.2.1
Reactivity of metal Complex - Inert and Labile Complexes.
6.2.2
Valence Bond and Crystal Field applications.
Kinetics of Octahedral Substitution
6.3.1
Nucleophilic Substitution
6.3.2
Hydrolysis Reactions
6.3.3
Factors affecting Acid Hydrolysis
6.3.4
Base- Hydrolysis-Conjugate Base Mechanism
6.3.5
Anation Reaction
6.3.6
Reactions without Metal-Ligand Bond-Cleavage
6.4
Let Us Sum Up
6.5
Check Your Progress: The Key
192
6.0
INTRODUCTION
Metal complexes are generally classified as 'Labile" and 'Inert' with
reference to their reactivity. The ability of a complex to engage itself in
reactions involving the replacement of one or more ligands in its
coordination sphere by other ligand is called lability of the complex. The
complexes that undergo rapid substitution are termed labile. Where as
those with low rates of substitution are called inert. However, the degree
of lability or inertness of a transition metal complex can be correlated
with the d-electron configuration of the metal ion. Nearly half of all
reactions of transition metal complexes may be considered substitution
reactions, while the remaining half are redox-reactions.
Equilibrium and kinetics play important and central part, for
determining the outcome of inorganic reactions. It is often helpful to
understand the mechanism of the reactions. A chemist who wishes to
synthesise an octahedral complex, as an example, must have some idea of
lability of the complex in order to choose appropriate experimental
conditions for synthesis.
Because the mechanism is rarely known finally and completely, the
nature of the evidence for a mechanism should always be kept in mind in
order to recognise what other possibilities might also be consistent with
it. In the first part of this unit we describe how reaction mechanisms are
classified, and distinguish between the steps by which the reaction takes
place and the details of the formation of the activated complex. Then
these concepts are used to describe the currently accepted mechanisms for
the substitution reactions of complexes.
However, you may recall what you have already studied about the
basic concept of kinetics and of current views on the nature of
193
substitution at a saturated carbon atom, as inevitably organic chemical
thinking has expected a great influence on the interpretation of the
kinetics of inorganic reactions.
6.1
OBJECTIVES
The main aim of this unit is to look in detail at the evidence and
experiments that are used in the analysis of reaction pathways and
develop a deeper understanding of the mechanism of substitution
reactions of d-block metal complexes. After going through this unit you
should be able to:

discuss the energy profile of a reaction and explain lability in
terms of VBT and CFT principles;

describe
nuelcophilic
substitution
reactions
in
octahedral
complexes in terms of SN1 and SN2 reaction mechanisms, and the
evidences supporting them;

explain acid- and base-hydrolysis reactions and their mechanisms;

explain water exchange (Anation) is a binuclear reaction and the
rate of this reaction depends upon the nature of the metal ion; and

6.2
describe the catalysts form octahedral substitution reactions;
ENERGY PROFILE OF A REACTION.
Why does a chemical reaction take place? What happens in a
chemical reaction? Answer of these and similar other questions are
important for a chemist; so that he can have control over a chemical
reaction and can either complete it or stop it, according to the need.
In order to convert reactants into product, it is necessary that the
groups or the atoms, linked in what ever manner, in the reactant
molecules should separate (may be partially) and then reunite (re-link) in
194
the form of the products. Unless this takes place using a suitable
mechanism, the reaction will not take place. On the thermodynamic basis,
the possibility of conversion of reactants, into products is only when the
state of disorder and the bond-energies in the products are relatively high.
Both of these, affect the direction of a chemical change and on the effect
of these depends the important thermodynamic functions Gibb's free
energy, G. For a chemical reaction, free energy is related with the heat
content or the useful energy, H , and the disorder, S , according to the
following relation:
G = H - T S
That is, a chemical reaction will go in the direction in which there
is decrease in free energy, i.e. G should be negative.
In order to a chemical reaction takes place, (i) the total bond
energies in the product are stronger then that in the reactants and the total
disorder (entropy) of the products is high or (ii) the total bond forces in
the products are stronger than that in the reactants and the products is
less, but T S is greater than H or (iii) the total bond forces in the
products are weaker as compared to that in the reactants, but the increase
in entropy is so high that it compensates the energy absorbed.
6.2.1 Reactivity of Metal Complex - Inert And Labile Complexes
Almost all the reactions of transition metal complexes may be
divided into two categories; (a) Substitution reactions and (b) Redoxreactions. In coordination chemistry rate of a reaction is equally
important, as the reaction equilibrium.
The ability of a complex to engage itself in reactions involving the
replacement of one or more ligands in its coordination sphere by other
195
ligands is called the lability of the complex. The complexes that undergo
rapid substitution (half time period or T1/2 or reaction rate k is used to
denote the speed of the reactions) are termed labile, whereas those with
low rates of substitution are called inert.
The inertness of the complex has nothing to do with the stability as
determined thermodynamically. Thus, [Ni(CN)4]2-,[Mn(CN)6]3- and
[Cr(CN)6]3- all have high stability constants. Yet the rate of exchange of
CN- by the labelled 14CN- gives half time period as 30 S, 1 h and 24 days,
respectively. Therefore [Ni(CN)4]2- is labile, while [Cr(CN)6]3- is inert.
Similarly [Fe(H2O)6]3+ (bond energy = 690 KJ mol-1) is labile while
[Cr(H2O)6]3+ is inert.
[Ni(CN)4]2- is thermodynamically stable but kinetically labile but
[Co(NH3)6]3+ is kinetically inert but thermodynamically unstable.
6.2.2 Valence Bond (VBT) And Crystal Field (CFT) Applications
(a)
VBT Application
According to VBT Octahedral Complex are of two types:
i. Outer-Orbital Complexes which involve sp3d2 hybridisation.
ii. Inner-Orbital Complexes which involve d2sp3 hybridisation.
The two d-orbitals involved in sp3d2 and d2sp3 hybridisation are
dx2-y2 and dz2 eg set orbitals.
1. Outer-Orbital Octahedral Complexes
Outer-orbital octahedral complexes (sp3d2 hybridisation) are
generally labile for example the octahedral complexes of
Mn2+(3d5)Fe2+, Fe3+(3d5) Co2+ (3d7) Ni2+ (3d8) Cu2+ (3d9) and Cr2+
(3d4) exchange ligands rapidly and hence are labile. This is
196
because, the use of outer d-orbitals does not make effective overlap
between metal and ligand orbitals resulting in weaker bonds.
2. Inner-Orbital Octahedral Complexes
The six d2sp3 hybrid orbitals are filled with the six electron pairs
denoted by the 6 ligands. dn electrons of the central metal will
occupy dxy, dyz and dxz orbitals. These complexes are inert as the
use of inner d-orbitals results in an effective overlap between metal
and ligand orbitals giving stronger bonds. Inner orbital octahedral
complexes are given in the Table 6.1 which explain the following
observations:
a. In the labile inner-orbital octahedral complexes there is at least
one d-orbital of t2g set empty, so that this empty d-orbital may
be used to accept the electron pair from the incoming ligand in
forming the transition state with coordination number
seven(unstable), which finally stabilise in to an octahedral
complex (Coordn. No. 6), removing one ligand (Fig 6.1).
Fig. 6.1
b. In the inert-orbital octahedral complexes every d-orbital of t2g
set (i.e. dxy, dyz and dxz) contains at least one electron, and
have no vacant orbtial to link an extra ligand.
197
Table 6.1 Distribution of dn-electrons in various t2g orbitals for
labile and inert inner-orbital octahedral complexes
(according to VBT)
n
Type of the
d
Distribution of dn electron (shown by
Example of central
complex
configu arrows) in t2g orbitals. Electrons shown by
metal ions
ration crosses in eg orbitals have been donated by
six ligands to enter d2sp3 hybrids and are
in opposite spins.
d
s
p
t2g
eg
2 2
xy yz zx x -y z2
px py pz
0
inner
d
xx xx xx xx xx xx Sc(+3), Y(+3), rare
orbital
earth (+3), Te(+4),
labile
Zr(+4), Hf(+4), Ce(+4),
octahedral
Th(+4), Nb(+5), Ta(+5),
complexes
Mo(+6), W(+6)

d1
xx
xx xx xx xx xx
Ti(+3), V(+4), Mo (+5),
W Re(+6)
d2


xx
xx xx xx xx xx
Ti(+2), V(+3), Nb (+3),
Ta(+3), W(+4), Re(+5),
Ru(+6)
inner
orbital
inert
octahedral
complexes
d3



xx
xx xx xx xx xx
V(+2), Cr(+3), Mo(+3),
W(+3), Mn(+4), Re(+4)
d4
 
d5



xx

xx xx xx xx xx
xx
[Cr(CN)6]4-, Mn(CN)6]1
Re(+3), Os(+3), Ir(+4)
4xx xx xx xx xx [Mn(CN)6] , Re(+2),
Fe(CN)6]3- Ru(+3),
Os(+3), Ir(+4)
d6



xx
4xx xx xx xx xx [Fe(CN)6] , Ru(+2),
Os(+2), Co(+3)
(except Co Fe34Rh (+3), Ir (+3)
198
Table 6.2 Loss in CFSE, E  (in the units of Dq) in the formation of
a pentagonal bipyramidal intermediate in octahedral substitution reactions
on the basis of SN2 associated mechanism
SN2 association mechanism Octahedral (oct.)Pentagonal bipyramidal
(pent.bipy.)
(C.N. = 6)
dn ion
(C.N. = 7)
Strong ligand field (spin-paired or
Weak ligand field (spin-paired or
low-spin complexes)
low-spin complexes)
Oct.
pent.bipy.
(C.N.= 6)
(C.N.= 7)
d0
0 Dq
0 Dq
d1
4
d2
E
E
Oct.
pent.bipy.
(C.N.= 6)
(C.N.= 7)
0 Dq
0 Dq
0 Dq
0 Dq
5.28
0
4
5.28
0
8
10.56
0
8
10.56
0
d3
12
17.74
-4.26
12
7.74
-4.26
d4
16
13.02
-2.98
6
4.93
-2.07
d5
20
18.30
-1.70
0
00
0
d6
24
15.48
-8.52
4
5.28
0
d7
18
12.66
-5.34
8
10.56
0
d8
12
7.74
-4.26
12
7.74
-4.26
d9
6
4.93
-1.07
6
4.93
-1.07
d10
0
0.00
0
0
0.00
0
Table 6.3 Loss in CFSE, E  (in the units of Dq) in the formation of a
pentagonal bipyramidal intermediate in octahedral substitution reactions
on the basis of SN1 associated mechanism
199
dn ion
d0
SN1 association mechanism
Octahedral (oct.) Syware Pyramidal (Squ. pyi)
(C.N. = 6)
(C.N. = 5)
Strong ligand field (spin-paired
Weak ligand field (spin-paired
or low-spin complexes)
or low-spin complexes)
Oct.
pent.bipy.
E
Oct.
pent.bipy.
E
(C.N.= 6) (C.N.=5)
(C.N.= 6) (C.N.= 5)
0 Dq
0 Dq
0 Dq
0 Dq
0 Dq
0 Dq
d1
4
4.57
0
4
4.57
0
d2
8
9.14
0
8
9.14
0
d3
12
10.00
-2-00
12
10.00
-2
d4
16
14.57
-1.43
6
9.14
0
d5
20
19.14
-0.86
0
00
0
d6
24
20.00
-4.00
4
4.57
0
d7
18
19.14
0
8
19.14
0
d8
12
10.00
-2.00
12
10.00
-2
d9
6
9.14
0
6
9.14
0
d10
0
0.00
0
0
0.00
0
The value of CFSE mentioned are in the units of Dq and have been
given for both the fields viz. strong field and weak field and for both the
mechanism (SN1, and SN2).
Negative values of E denotes a loss of CFSE when octahedral
complex is changed into an activated complex which may be square
pyramidal or pentagonal bipyramidal. If the CFSE of the activated
complex is greater than that of octahedral complex. E  has been given
zero value which shows that these complexes do not loose CFSE when
they are changed into activated complexes.
200
The octahedral complexes formed by the ions for which there is
large loss in CFSE are least labile i.e. such complexes are inert.
On the other hand octahedral complexes given by ions for which
there is little or no loss in CFSE are labile i.e. such complexes react
rapidly. Thus we see:
i Both high spin and low spin octahedral complexes of d0, d1 and d2
ions will react rapidly, i.e. these are labile complexes, in which
there is no loss in CFSE.
ii. According to VBT inner-orbital octahedral complexes of d3, d4, d5,
and d6 ion are inert while these are called low spin or spin paired
complexes according to CFT.
CFT predicts that low spin complexes of these ions are also inert
whether the mechanism is assumed to be SN1 or SN2 in which
CFSE values decreases.
The ion with maximum loss of CFSE will form the most inert
complex. Thus the order of inertness of low spin complexes formed by d 3,
d4, d5 and d6 ions is:
Order of inertness
:
d6 > d 6 > d4 > d5
Loss of CFSE for SN1 mechanism
:
-4.00>-2.00>-1.43>-0.86
Loss of CFSE for SN2 mechanism
:
-8.52-4.26-2.98-1.70
The order of reactivity will be reverse of the above i.e. the order of
reactivity will be d6 > d3> d4 > d5 it is supported by the following facts:
i. High spin octahedral complexes of d3 ion will react slowly, i.e. these
are inert complexes because for this ion there is substantial loss in
CFSE whether the substitution mechanism is SN1 or SN2.
201
ii. High spin octahedral complexes of d5 ion react rapidly i.e. these are
labile complexes, since there is no loss in CFSE.
iii. Both high spin and low spin octahedral complexes of d8 ion are inert.
According to VBT d8 ion [3dxy2, 3yz2, 3dxz2, 3d(x2-y2),
3dz2] will form outer orbital complexes and will be labile. Therefore
in case d8 ion VBT & CFT gives different predictions.
iv. Both high spin and low spin octahedral complexes of d 10 ion are
labile.
Factors Affecting the Liability of Complex
1. Charge of the metal ion: For the isoelectronics complexes there is
a decrease in lability with the increase of the charge of the central
metal ion.
i. The order of lability of the complex is as follows:
Lability order : [AlF6]3- > [SiF6]2 > [PF6]- > [SF6]0
Cationic charge : +3
<
+4
<
+5
<
+6
ii. The rate of water exchange represented by:
[M(H2O)6]n + 6H2O*  [M(H2O*)6]n+ + 6H2O
decreases with the increase of cationic charge in the series
Rate of water exchange:
[Na(H2O)6]+ > [Mg(H2O)n]2+ > [Al(H2O)6]3+
Cationic Charge +1
<
+2
<
+3
2. Radii of the Central ion : Complexes having central atoms with
small ionic radii react more slowly than those having larger central
ions i.e. the lability increase with the increase of ionic radius e.g.
Order of liability: [Mg(H2O)6]2+<[Ca(H2O)6]2+< [Sr(H2O)6]2+
Cationic Size (A)
0.65 <
202
0.99 <
1.13
3. Charge to Radius Ratio Values: Octahedral complexes having
the central metal ion with the largest charges to radius ratio will
react slowest (Fig. 6.2).
i. The first row transition elements [Ni(H2O)6]2+ (a d8 system)
has the largest value of half life i.e. it reacts slowest. The
hydrated M2+ ions [M(H2O)x]2+ of the first row transition
elements are all high spin complexes.
ii. [Cu(H2O)6]2+ reacts most rapidly because the 2 water
molecules above and below the square plane of the tetragonal
distorted octahedral shape of [Cu(H2O)6]2+ are exchanged. The
remaining four H2O molecules lying in the square plane react
slowly.
Fig. 6.2: Half-lives (in sec) at 25oC for the exchange
of water by some hydrated metal ions.
4. Geometry of the Complex: Four co-ordinated complexes react
more rapidly than analogues 6-co-ordinated complexes e.g. the
very stable [Ni(CN)4]2+ undergoes rapid exchange with 14CN-,
[Ni(CN)4]2+ + 414CN-  [Ni(14CH)4]2- + 4CN
203
while 6-co-ordinated complexes like [Mn(CN)6]4- and [Co(CN)6]3have the same stability as [Ni(CN)4]2+. The greater reactivity of
4-co-ordinated complexes may be due to the fact there is enough
room round the central ion for the entry of a 5 th group into the coordination sphere to form on activated complex.
Check Your Progress-1
Notes : (i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
A.(i)For a reaction to go in the forward direction G should
be................................
(ii)According to the thermodynamic relation G = .........................
That is for conversion of the reactants into the products, the
bond energies and the state of disorder should be....................i.e.
the value of the  H should be.....................and that of T  S
should be...............................
B.(i) According to VBT, generally labile complexes are........................
complexes, while the inert complexes are.................complexes.
(ii) Inner orbital complexes may be labile, if they have at
least.............. in.........................set is vacant, e.g. in..................
(iii) According to CFT inert complexes have............................values
of ................................
204
6.3
KINETICS OF OCTAHEDRAL SUBSTITUTION
Substitution reactions involve the activated complex which is most
unstable changes to give the product x-y and z. Thus the various steps
responsible for the reaction are
X + Y - Z  X.......... Y..........Z  X - Y + Z
Reactants
Activated complex
(Transition State)
Unstable
Products
The difference in energy between the reactants and the activated
complex is called activation energy.
These reaction involves two process (1) SN1 and SN2
1. In SN1 process the rate-determining slow step is a metal-ligand
bond breaking step, since the co-ordination No. of the complex
MX5Y (=6) is decreased to 5 which is the co-ordination number of
the intermediate MX5.
For a ligand replacement reaction of the general type
[LnMX] + Y = [LnMY] + X
(For simplicity all charges are omitted), the mechanism analogous
to unimolecular nucleophilic substitution (sN1) at a carbon atom
would be:
slow
 [LnM] + X
[LnMX] 
fast
 [LnMY] + X
[LnM] + y 
The rate of SN1 mechanism is first order with respect to
MX5Y, i.e. the rate-determining step in this mechanism is
unimolecular.
On the other hand the rate determining step for SN2
mechanism is bimolecular, i.e. its rate of reaction is second order
205
first order with respect to MX5Y and first order with respect to Z.
Thus
for SN1 mechanism rate = K [MX5Y],
and
for SN2 mechanism rate = K [MX5Y][Z]
Here, it may be mentioned that the kinetic data would be
equally compatible with ion-pair formation (if both reactants are
ions) followed by a unimolecular reaction of the ion-pair:
k1
[Ln MX] + Y 
[Ln MX] Y
slow
[Ln MX] y 
 [Ln MY] + X
This leads to
k k [ L MX ][Y ]
d
[Ln MY] = 1 2 n
dt
k 1  k 2
= k[Ln MX][Y]
where
k=
k1k 2
k 1  k 2
Detailed investigation of such a reaction can lead to a value for
k1/k-1, the equilibrium constant for ion-pair formation.
6.3.1 Nucleophilic Substitution
As has been motioned, nucleophile substitution reactions in
octahedral complexes follow either of the two mechanisms, the
dissociation mechanism or the SN-1 mechanism and the association
mechanism or the SN-2 mechanisms. The rate determining step in
association or dissociation, may be worked out by analysing the rate-laws
of the reactions taking place and the specific conditions under which the
reactions take place. The difference in these two mechanisms depends on,
206
whether the rate determining steps is the formation of a new Y...............M
bond or the dissociation of an old M...................X bond.
(a) SN-1 or Dissociation Mechanism
The nucleophilic substitution unimolecular reaction actually
proceeds in two steps. In the first, slow and rate determining step,
one ligand Y is lost and a five coordinated intermediate is formed.
In the second step the short-lived penta-coordinated
intermediate of very limited stability is attacked rapidly by the
nucleophilic reagent, Z to give the complex, MX5Z.
There two steps are diagrammatically shown in Fig. 6.3
Fig. 6.3: SN1 or dissociation mechanism for the substitution
reaction MX5Y + Z  MX5Z + Y
For the SN1 mechanism, the following points are important:
(1) The trans effect of the ligands would not be operative due to the
dissociation of the ligand completely from the octahedral
complex.
(2) The rates of SN1 substitution (k1) should be inversely proportional
to the strength of the Co-L bond, and depend on the charge, steric
factors, and chelating effects of the leaving group L.
207
(3) Increase in the electron density on Co atom by the electron
donors in SNn should assist the M-L bond breaking.
(4) k1 is independent of the nature of E as well as its concentration
except for the OH- group for which the reaction is of the second
order.
(5) Cis effect. Ligands having another pair of electrons like CNS - or
OH- increase the rate of hydrolysis of the complexes about ten
fold when present cis to L, as compared to the rate when they are
present trans to L. This is due to the stabilization of the square
pyramidal complex by the electron pair donation by OH- or CNSalong the Cis position through p-d-  bonding (Fig. 6.4). No
rearrangement takes place and the product is 100 percent Cis
isomer. The ligands that do not show the eis effects are those that
do not have an extra pair of electrons (NH3) or are themselves 
acceptors (NO2-, CO, NO, etc.).
Fig. 6.4: Cis-effect
From Table 6.4, it can be seen that for the formation of the
5-coordinate intermediate, high energy changes are required for
the low spin d3, d6 and d8 and high spin d3 and d8 ions. Hence,
these complexes do not favour the SN1 mechanism for the
substitution.
Table 6.4 Changes in LFSE (in Dq) for Changing a 6coordinate Complex to a 5-Coordinate (SP) or a 7-Coordinate
(Pentagonal Bipyramid) species.
208
System
do, d10
High spin
CN = 5
CN = 7
0.00
0.00
Low spin
CN = 5
CN = 7
0.00
0.00
d1
0.57
1.28
0.57
1.28
d2
1.14
2.56
1.14
2.56
d3
-2.00
-4.26
-2.00
-4.26
d4
3.14
-1.07
-1.43
-2.98
d5
0
0
-0.86
-1.70
d6
0.57
1.28
-4.00
-3.52
d7
1.14
2.56
1.14
-5.34
d8
-2.00
-4.26
-2.00
-4.26
d9
3.14
-1.07
3.14
-1.07
+ value indicate gain in CFSE while - values indicate loss in
CFSE.
(b) SN-2 or Association Mechanism
SN-2 or the nucleophilic bimolecular substitution reaction
also proceeds through two steps:
The first step is slow step and involves the attachment of the
incoming nuclepohile, Z to MX5Y to form a seven-coordinate
unstable intermediate (perhaps transition state) which is probably
pentogonal bipyramidal in shape. Obviously it is a metal-ligand
bond-making step.
( Z )


MX5Y Slow
(C.N.=6)
MX5YZ
Unstable seven-coordinatee
Intermediate (C.N.=7)
This reaction is rate-determining and bimolecular because
two reactants viz MX5Y and Z are involved in this step. Thus the
rate of this rate-determining reaction is of second order: first order
209
with respect to the complex, MX5Y and first order with respect to the
entering ligand, Z, i.e.,
Rate of reaction = K[MX5Y][Z]
In the second step either at the same as Z adds to MX5Y or
shortly thereafter, Y leaves MX5YZ rapidly to give MX5Z. This is a
fast step.
MX5YZ
Unstable seven-
Fast


MX5Z
-Y
(C.N.=6)
coordinatee Intermediate
(C.N.=7)
Both these steps are shown diagrammatically in Fig. 6.1
This mechanism is similar to Eigen-Wilkins Mechanism,
which presents formation of the association complex [L-MX5-Z] in
the pre-equilibrium step: Thus the following equilibrium will be
established:
LMX5 + Z  MX5. Z; K =
[ LMX 5 .Z ]
[ LMX 5 ][ Z ]
The value of the equilibrium constant, K, for the association
complex, may be obtained using Fuoss-Eagan equation,
K=
where,
4
3
-v/RT
 a NAe
3
a = Nearest reach-distance
v = Coulomb potential energy at a-distance
NA = Avogadro number = (Z1Z2e2/4  ea)
As in the octahedral complexes, the six ligands are already
present along the three C4 axes along which the eg orbitals are
concentrated, the t2g orbitals that lie along the C2 axes most probably
have to be approached by the seventh ligand to form the associated
complexes in the SN2 process.
210
Hence, if the t2g orbitals are filled (Co2+ in low spin
octahedral complexes), the higher activation energy required to
empty one of the t2g orbitals will make the complex inert.
Table 6.4 also shows that due to the loss of the CFSE
energies, the d3 and low spine d6 ions require highest activation
energies, followed by d7, d8 (Ni2+complexes are labile due to the
expulsion of ligand by the eg orbitals) and high spin d3 and d8 ions.
Thus, SN-1 and SN-2 reactions differ in the following points:
(i) In SN1 process the rate-determing slow step is a metal-ligand
bond breaking step, since the coordination number of the
complex, MX5Y (=6) is decreased to 5 which is the coordination
number of the intermediate, MX5. On the other hand in SN2
process the rate-determining step involves a metal-ligand bond
making step, since C.N.=6 is increased to 7.
(ii) The rate of SN1 mechanism is first order with respect to MX5Y,
i.e., the rate-determining step in this mechanism is unimolecular.
On the other hand the rate-determining step for SN2 mechanism is
bimolecular, i.e. its rate of reaction is second order: first order
with respect to MX5Y and first order with respect to Z. Thus:
for SN1 mechanism rate = K[MX5Y]
and for SN2 mechanism rate = K[MX5Y][Z]
6.3.2 Hydrolysis Reactions
The substitution reactions in which a ligand is replaced by a H 2O
molecule or by OH- groups are called hydrolysis reactions. They are of
two types (a) when an aqua complex is formed by the replacement of a
ligand by H2O molecules are called acid hydrolysis or equation reactions,
211
while (b) the reactions, in which a hydroxo complex is formed by the
replacement of a ligand by OH- group are called base hydrolysis reactions.
Acid hydrolysis reactions occur in neutral and acidic solutions
(pH<3) while base hydrolysis reactions occur in basic solution (pH  10).
Examples are:[CoIII(NH3)5Cl]2+ + H2O[CoIII(NH3)5(H2O)]3+ + Cl- Acid hydrolysis
[CoIII(en)2ACl]+ + H2O  [CoIII(en)2A(H2O)]2+ + Cl- reaction
[A = OH-, Cl-, NC-, NO2-]
[Co(NH3)5Cl]2++OH-[Co(NH3)5(OH)]2++Cl- (Base hydrolysis reaction)
(a) Acid Hydrolysis or Aquation :
When NH3 or ammines like ethylene diamine or its
derivatives co-ordinated Co(III) are replaced very slowly by H2O
molecules and hence in acid hydrolysis only the replacement of
ligands other than amines is usually considered.
The rate of hydrolysis of the reaction is of first order.
[Co(NH3)5X]2+ + H2O[Co(NH3)5(H2O)]3+ + XThe rate of hydrolysis reaction is of first order.
In aqueous solution the concentration of water is always
constant, the effect of changes in water concentration on the rate of
the reaction cannot be determined.
The rate law K = K1[Co(NH3)5X]2+[55.5] does not indicate
whether these reactions proceed by an SN2 displacement of X by
H2O or by an SN1 dissociation followed by the addition of H2O.
212
The rate law for acid hydrolysis at low pH thus becomes
(If
X
d
[Co(NH3)5X] = kA[Co(NH3)5X]
dt
is
the
anion
of
a
weak
acid,
a
term
kH+[Co(NH3)5X][H+] is added.) As we have shown previously, such
a rate law is compatible with either a slow dissociation of the
complex into [Co(NH3)5]3+ and X or replacement of X by H2O as the
rate-determining step. In order to try to decides between these
alternatives, the rates of hydrolysis of a series of complexes of
formula [Co(AA)2Cl2]+, where AA is a substituted ethylendiamine,
were examined. For replacement of a single chloride ion at pH 1 the
order found for values of kA was
CH2NH3
CH3CHNH2
<
CH2NH3
CH3CHNH3
<
CH2NH2
(CH3)2CNH2
<
CH3CHNH2
(CH3)2 CNH2
Such an acceleration of substitution by bulky ligands
suggests that the dissociative mechanism is operative; although
introduction of methyl groups must have some inductive effect, the
variation in base strengths among the diamines is very much less
than the variation in rate constants for the hydrolysis of their cobalt
(III) complexes, and it seems reasonable to attribute the kinetic
effect mainly to steric factors. Now since steric factors favour S N1
reactions, this is evidence for the dissociative mechanism. Further
evidence for this mechanism is provided by:
(a) a general inverse correlation between the rate of replacement of
X in [Co(NH3)5X] and the formation constant of the
[Co(NH3)5X] complex from [Co(NH3)5(H2O)]3+ and X, and
213
(b) the decrease in the rate of the exchange reaction
[Co(NH3)5(H2O)]3+ H218O = [Co(NH3)5(H218O)]3+ + H2O at high
pressures.
6.3.3 Factors Affecting Acid Hydrolysis
(i)
Effect of Charge on the Complex:
The value of rates of acid hydrolysis of some Co(III) complexes at
pH=1 shows that the divalent monochloro complexes react about 100
times slower than the monovalent dichloro complexes.
As the charge of the complex increases, a decrease in rate is
observed and the acid hydrolysis of the divalent complexes like
[Co(NH3)4(H2O)Cl2]2+ occurs in two steps:
[Co(NH3)4 (H2O)Cl]2+ +
slow

 Cl
6-co-ordinate complex
[Co(NH3)4(H2O)]3+ +
[Co(NH3)4(H2O)]3+ + Cl5-co-ordinate Intermediate
fast

 H 2O
[Co(NH3)4(H2O)2]3+
The acid hydrolysis represented by equation (1) would proceed
more rapidly than that represented by equation (2) because the separation
of a negative charge in the form of Cl ion from a complex ion with higher
charge is more difficult.
(ii)
Effect of Chelation
When NH3 molecules in [Co(NH3)5Cl]2+ complex ion are replaced
partially or completely by polyamines like en, trien, diene, tetraene etc,
the rates of the reaction of the divalent complex ions shows that as the
number of -CH2-CH2 or -(CH2)2-chelated links increases the rate values
decreases.
214
The replacement of NH3 molecules by polyamines increases the
size of the complex i.e. the chelated complex has larger size. The larger
the size of the ion less its solvation energy will be and hence less easily it
will be formed. Thus the stability of the transition state in which the Cl
ion is only partially lost and in which the solvation is less efficient will be
reduced. The rate of equation is slowed down by chelation because of
reduced stability of the transition state due to less efficient solvation.
(iii)
Effect of Substitution on ethylene diamine
When H atoms on carbon atom or on nitrogen atom of en groups of
trans [Co(en)2Cl2]+ are replaced by the alkyl groups like CH3,C2H5 etc.
the ligand becomes more bulky, if the strained complex having bulky
ligand reacts by SN1, dissociative mechanism and co-ordination number 6
is reduced into 5 co-ordinated intermediate, on the other hand if the
strained complex reacts by SN2 displacement process, the crowding on
the complex is increased as it is converted into a transition state of coordination number seven. The rate of hydrolysis of trans [Co(AA2 Cl2)]+
at 25oC and pH=1 corresponding to the replacement of only one Cl- ion
by H2O molecule are given. Here AA is the diamine.
(iv)
Effect of Leaving Group
The rate of reaction of [Co(NH3)5X]2+ corresponding to the
replacement of X with H2O molecule depends on the nature of X because
the bond breaking step is important in rate determining step. The
reactivity of X-groups decreases in the order.
HCO3->NO3->I->Br->Cl->SO4-->F->CH3COO->SCN-<NO2
6.3.4 Base- Hydrolysis-Conjugate Base Mechanism
The base hydrolysis reaction represented by following equation:
[Co(NH3)5Cl]2++ OH-  [Co(NH3)5(OH)]2++ ClIt involves following two mechanisms.
215
1.
SN-2, Displacement Mechanism
The reaction proceeds as:
+
2+
(  OH )
fast

 [Co(NH3)5(OH)Cl] 
 [Co(NH3)5(OH)] + Cl
[Co(NH3)5Cl]2+ slow
(C.N.=6)
(C.N.=7)
Rate of Reaction =
=
The rate law is 2.
(C.N.=6)
K[Complex][base]
K[Co(NH3)5Cl][OH-]
d
[Co(NH3)5Cl] = KB[Co(NH3)5Cl][OH-]
dt
SN-1 Displacement Mechanism:
The complex which acts as a Bronsted acid is converted into
its conjugate base (CB), [Co(NH3)4(NH2)Cl] + which is obtained by
removing a proton H+ from the amino group present in the
complex. CB is an amido complex, since it contains an amido
group.
SN-1 mechanism fails to explain quite a few observations:
(1) 7-coordinate complexes are not very stable.
(2) The value of kn is nearly 104 times higher than kA. Why
should hydroxyl ions posses the exceptionally high
nucleophilic activity as compared to the similar anions?
(3) If a proton cannot be abstracted from N5 ligands (e.g.,
[Co(py)4-Cl2]+ or [Co(CN)5Cl]3-), reaction rate for hydrolysis
is very low.
To overcome the above difficulties, an alternative
mechanism is proposed by Garrick (1987). In this case the
OH- ions abstract a proton form a ligand in N5 group giving
CB of the ligand. This undergoes the dissociative
mechanism as shown below:
+
fast
 [(NH3)4Co(NH2)Cl] +H2O (6.1)
[(NH3)5CoCl]2++OH- 
2+

 [(NH3)4Co(NH2)] + Cl
[(NH3)4Co(NH2)Cl]+ slow
(6.2)
2+
2+
fast
 [(NH3)5Co(OH)]
[(NH3)4Co(NH2)] +H2O 
(6.3)
216
The rate determining step is the dissociation of the
amido complex given in Eq.(6.2) whose concentration would
depend upon the concentration of hydroxyl ions present. This
is the SN1CB process.
The rate law will bed
= [Co(NH3)5OH] =
dt
K1K 2[Co ( NH 3 )5 Cl ][OH ]
K 1[ H 2O] 2 K 2[ H 2O]
= K[Co(NH3)5Cl][OH]
where,
K=
K1 K 2
K 1[ H 2O]2  K 2 [ H 2O]
Though it seems very unlikely that reduction in one
positive charge form [Co(NH3)6]3+ to [Co(NH3)5(NH2)]2+
should increase the reaction rate enormously, it is possible
that through a  bonding intermediate, the stability of the 5coordinate complex is increased (Fig. 6.5).
Fig. 6.5: Stabilization of the intermediate 5-coordinate species
through the resonance effects involving NH2 group.
The SN1 CB mechanism does not explain the
following observations. (i) The conjugate base readily
dissociates and releases the ligand L; and (ii) the
concentration of the conjugate base is very low due to the
basic nature of the ligands, and should be present only as a
very small fraction of the concentration of the complex
present.
217
Direct and Indirect Evidences in Favour of Conjugate Mechanism:
Equation 6.1 requires that the reacting complex should have at
least one Photonic hydrogen atom (H+) on a non-leaving ligand so that H+
may transfer to OH- to form its conjugate acid H2O and conjugate base,
[Co(NH3)4(NH2)Cl]+ of [Co(NH3)5Cl]2+ which acts as an acid. Thus a
complex having no proton should react with OH- much more slowly and
the rate of reaction would be independent of the concentration of OH -. It
is observed that the complexes like [Co(Cn)2Br] and trans [Co-(Py)4Cl2]+
which does not have N-H hydrogen undergo hydrolysis much more
slowly in basic solution at a rate which is independent of [OH-] over a
wide range. Thus in the absence of an acidic portion on the ligands an
SN1 CB mechanism is not possible.
Such complexes undergo rapid base hydrolysis supports the SN1
CB mechanism and the acid-base properties of the complexes are more
important to the rate of reaction, than the nucleophilic properties of OH.
Thus both the mechanisms give the same rate laws and the same
hydroxo products in aqueous solution, because water is a good cocoordinating agent. The rate of formation of [Co(en)2(NO2)Y]+ depends
only on the concentration of the base, OH, not on the nature or
concentration of Y-, OH- and piper; dine are used as catalysts while N3-,
NO2-, SCN- ion are used as nucleophiles.
In SN1CB mechanism the reactions of [Co(NH3)5Cl]2+ and OH- in
aqueous solution at 25oC in presence of H2O, when H2O2 is added to the
reaction mixture of [Co(NH3)5Cl]2+ and OH-, the reaction between OHand H2O2 occurs as:
OH- + H2O2  H2O + HO2-
218
Which increase the rate of base hydrolysis reaction and form
peroxo products.
On the other hand if the reaction occurs by an SN1CB mechanism
the addition of H2O2 to the reaction mixture should reduce the rate of
base hydrolysis reaction compared to OH- because of the reduction in the
concentration of OH- ions. The rate of SN1CB reaction is directly
proportional to the concentration of OH-.
6.3.5 Anation Reaction
The reaction in which an aquo ligand (i.e. H2O molecule) from an
aquo complex is replaced from the co-ordination shell by some axion.
[Co(NH3)5(H2O)]3+ + X- [Co(NH3)5X]2+ + H2O
Thus we find that an anation reaction is the reverse of acid
hydrolysis reaction.
It shows that these are bimolecular reactions with a rate which
depends on the concentration of the complex and X. The same second
order kinetics would be observed for a unimolecular process.
3+
2+
slow
fast
 [Co(NH3)5] 
 [Co(NH3)5X] + H2O
[Co(NH3)5(H2O)]3+ 
Let us consider replacement of water in a species containing five
non-labile ligands such as [Co(NH3)5(H2O)]3+ , and let us reverse the
experimental procedure and attempt to infer kinetic behaviour from a
postulated mechanism. This is
k
[L5M(H2O] 
[L5M] + H2O
1
k
[L5M] + Y 
[L5MY]
1
Since Y competes with solvent water for the active intermediate
[L5M], the rate of formation of [L5MY] can be dependent on the
concentration of Y. On the other hand, there should be some high
219
concentration of Y at which the rate of replacement of water no longer
depends on the concentration of Y. The rate of formation of [L 5MY] at
this concentration should be equal to the rate of formation of [L5M] and
also equal to the rate of exchange of water between [L5M(H2O)] and the
solvent. Thus the rate of formation of [L5M] is given by
d
[L5M] = k1[L5M(H2O)] - k-1[L5M][H2O] - k2 [L5M][Y]
dt
According to the steady-state approximation, the concentration of
the very reactive [L5M] remains small and constant during the reactions
(i.e.
d
[L5M] = 0 at the steady state). Thus,
dt
[L5M] =
K1[[L5 M(H 2O]
K 1[ H 2O]  K 2 [Y ]
and
K K [L M(H 2O][Y ]
d
[L5MY] = 1 2 5
dt
K 1[ H 2O]  K 2 [Y ]
if k-1 [H2O] > k2[Y]
KK
d
[L5MY] = 1 2 [L5M(H 2O][Y ]
dt
K 1
and a second-order reaction will be observed. On the other hand,
if k2[Y] > k-1[H2O]
d
[L5MY] = k1 [L5M(H 2O]
dt
giving first-order kinetics with the overall first-order constant
equal to that for the dissociations of the aquo complex.
6.3.6 Reactions without Metal-Ligand Bond-Cleavage
Many a times, replacement of ligand takes place without breaking
a metal-ligand bond. Important examples of this fact are formation of
aquo-complex,
[Co(NH3)5H2O]3+,
220
from
carbon
a
to
complex,
[Co(NH3)5CO3] and nitrito complex [Co(NH3)5ONO]2+ from hydroxo or
aquo-complex, [Co(NH3)5H2O]3+.
The most likely path for the equation of the carbonato complexes
seems to be the electrophilic attack by the proton on the O atom bonded
to the metal, so that no O is found in the complex when the equation is
carried out in presence of H2O (Fig. 6.6).
Similarly the reaction of pentamineaquacobalt (III) with NO2- ion
is explained by the sequence in Fig. 6.7.
Fig. 6.6 Mechanism of substitution of carbonate group by water
through electrophilic attack by H2O+.
Fig. 6.7 Probable mechanism of substitution of [(NH3)5Co(OH)]2+ by
nitrite NO2- through an electrophilic attack.
221
Check Your Progress-2
Notes : (i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the unit.
(a) SN-1 and SN-2 mechanism of substitution reactions differ in (i) In SN-1 rate determining step is.............................process, while
that in SN-2 is .......................process.
(ii) The rate determining step in SN-1 is...................molecular, while
that in SN-2 is ....................... molecular.
(iii) In SN-1 rate = ......................................
while that in SN-2 rate = ......................................
(b)(i) For acid hydrolysis at low pH, the rate Law is ...................................=.................................................
(ii) The rate law of base hydrolysis reactions of an octahedral
ammine complex, by SN-1 CB process is d
= [Co(NH3)5OH] = .................................
dt
= ...................................
(iii) Formation of aquo-complex from a carbonato complex is an
example of substitution...................... bond, and involves..............
attack on....................
6.4
LET US SUM UP

In order to convert reactions in to products it is necessary that the
groups or the atoms linked in what ever manner in the reactants
molecules should separate and then reunite in the form of the
products.
222

On the thermodynamic basis, Gibbs free energy for the reaction
should decrease, in order to the reaction takes place, i.e.  G
should be negative.

Since,
 G =  H - T  S, hence the possibility of conversion of reactants
in to products is only when the state of disorder, and the bond
energies in the products, are relatively high;  H should be
negative and  S should be positive.

Complexes are generally classified as labile and inert with
reference to their reactivity. The ability of a complex to engage
itself in the reactions involving the replacement of one or more
ligands in the coordination sphere by other ligand is called lability
of the complex.

The complexes that undergo rapid substitution are termed labile;
where as these with law rates of substitution are called inert.

According to VBT, the inner orbital complexes (using d2sp3
hybridisation for octahedral complication) are inert while the outer
orbital complexes (using sp3d2 hybridisation) are labile.

The inner orbital complexes may be labile only when they have at
least one d-orbital in t2g set is vacant; e.g. [V(NH3)6]3+ ion,

According to CFT, complexes with high values of CFSE are inert,
while those with small values of CFSE are labile.

Substitutions of ligands generally follow one of the two, S N-1 or
SN-2 mechanisms.
223

In SN-1 or dissociation mechanism, the rate determining slow step
is a metal-ligand bond breaking step, since the coordination
number of the complex, MX5Y (=6) is decreased to 5 in the
intermediate, MX5, complexes:
Y
Z
[MX5Y] 
[MX5] 
[MX5Z]
(C.N.=6)
(C.N.=5)
(C.N.=6)
Thus, the rate of SN-1 mechanism is first order with respect
to MX5Y, i.e. the rate determining step is unimolecular.

In SN-2 process the rate determining step involves a metal-ligand
bond making step, with the increase of coordination number from
6 to 7:
Z
Y
[LMX5] 
[L-MX5Z] 
[MX5Z]
(C.N.=6)
(C.N.=7)
(C.N.=6)
Thus, the rate determine step in SN-2 process in bimolecular i.e. its
rate of reactions is second order; first order with respect to [MX 5L] and
first order with respect to Z.

For SN-1 mechanism, rate = k [MX5L]
For SN-2 mechanism, rate = k [MX5L][Z]

The substitution reaction in which a ligand is replaced by a H2O
molecule or by OH- group is known as hydrolysis reaction.
The reaction is called 'acid hydrolysis' or 'aquation' when an aquo
complex is formed by the replacement of a ligand by H2O molecule while
the reaction in which a hydroxo complex is formed by the replacement of
a ligand by -OH group is called base-hydrolysis. Acid hydrolysis occur,
in neutral and acid solutions (pH<3), while base hydrolysis occurs in
basic solutions (pH>10).
224

The rate of acid hydrolysis reaction is of first order [Co (NH3)5X]2+ + H2O  [Co(NH3)5(H2O)]3+ + XAs in oqueous solutions the concentration of water is always
constant, the rate law, K = K1 [Co (NH3)5X]2+ [55.5] does not indicate
whether these reactions proceed by and SN-2 displacement of X by H2O
or by SN-1 dissociation followed by the addition of H2O.

For base hydrolysis;
[Co (NH3)5Cl]2+ + -OH  [Co(NH3)5(OH)]2+ + ClSN-2 mechanism gives rate of the reaction = K[Co(NH3)5Cl]2+
[OH-] and the rate law 
d
[Co(NH3)5Cl] = KB[Co(NH3)5Cl][OH-].
dt
Gerick proposed SN-1CB mechanism for base hydrolysis reaction.
In this, the -OH ions abstract a proton from a ligand in N5 group,
giving the conjugate base of the ligand. This under goes the
dissociative mechanism:
+
Fast
[(NH3)5CoCl]2+ + OH- 
 [(NH3)4Co(NH2Cl)] + H2O
2+
[(NH3)4Co(NH2)Cl]+ Slow

 [(NH3)4Co(NH2)] + Cl
2+
Fast
 [(NH3)5Co(OH)]
[(NH3)4Co(NH2)]2+ + H2O 
Thus, the rate determining step is the dissociation of the amido
group. The rate law will bed
[Co(NH3)5OH] =
dt
K1K 2 [Co(NH3 )5 Cl][OH- ]
K 1[ H 2O] 2 K 2[ H 2O]
= K [Co(NH3)5Cl][OH-]

The reactions involving removal of coordinated water molecule are
known as 'anation' reactions:
[[Co(NH3)5(H2O)]3+ + X-  [Co(NH3)5X]2+ + H2O
225
This reaction is reverse of acid hydrolysis reaction. The same
second order kinetics would be observed for a unimolecular process:
3+
 [Co(NH3)5]
[Co(NH3)5(H2O)]3+  Slow
H 2O

Fast

 X
[Co(NH3)5X]2+
Many a times replacement of ligand takes place without breaking a
metal-ligand bond, e.g. formation an aquo-complex from a
carbonato complex. These involve the electrophilic attack by the
proton on the O-atom bonded to the metal.
6.5
CHECK YOUR PROGRESS: THE KEY
1(a)(i)  G should be negative.
(ii)  G =  H - T  S
That is .......................should be very high i.e.  H should be
negative and T  S should be positive.
(b)(i) Labile complexes are outer orbital complexes....................inert
complexes are inner orbital complexes.
(ii) One orbital in t2g set:
e.g. [V(NH3)6]3+
(iii) Have high values of CFSE.
2.(a)(i) Is metal-ligand bond breaking process that in SN-2 is metalligand bond making process.
(ii) SN-1 is Unimolecular.
SN-2 is bimolecular.
(iii) SN-1, rate = K[MX5Y]
SN-2, rate = K[MX5Y][Z]
226
(b)(i) Rate law is -
d
[Co(NH3)5X] =KA[Co(NH3)5X]
dt
K1 K 2 [Co(NH3 ) 5 Cl][OH - ]
(ii) =
K 1[ H 2 O] 2  K 2 [ H 2 O]
= K [Co(NH3 )5 Cl][OH- ]
(iii) Without breaking M-L bond.
Involves electrophilic attack of proton on oxygen bonded
with metal.
227
M.Sc. (Previous) Chemistry
Paper – I :
INORGANIC CHEMISTRY
BLOCK – III
UNIT – 7 : Reaction Mechanism of Transition Metal
Complexes-II
UNIT – 8 : Metal Ligand Equilibria in Solution
UNIT – 9 : Metal Clusters
UNIT – 10 : Isopoly and Heteropoly Acids and Salts.
Author – Dr. Purushottam B. Chakrawarti
Edtor – Dr. M.P. Agnihotri
228
UNIT 7
REACTION MECHANISM OF TRANSITION METAL
COMPLEXES-II
Structure
7.0
Introduction
7.1
Objectives
7.2
Redox – Reactions
7.2.1 Mechanism of one electron transfer reaction
7.2.2 Outer Mechanism.
7.3
Theories of Redox – reactions.
7.4
Potential Energy Diagrams
7.5
Marcus Theory
7.6
Factors Affecting Electron Transfer Reaction Rate
7.7
Inner Sphere Type Reactions
7.8
Let Us Sum Up
7.9
Check Your Progress : They Key
229
7.0
INTRODUCTION
In this unit we discuss the kinetics and mechanisms of redox
reaction for octahedral complexes. During redox reactions, are often very
fast. However, in recent years, the scope of schemical kinetics has been
much increased by developments in the study of fast reactions. Various
physico chemical methods are used to study these reactions; such as
isotopic tracers or magnetic resonance method.
The factors that influence the rates of electron transfer reaction are
many more. Transfer of electron from one species to other may take outer
sphere root or the inner sphere one. During the first root, bond-formation
and bond cleavage does not take place; while during the later root excited
species are involved. However, the quantitative understanding of the rates
of inorganic reactions is far from secure, and in most cases all that it is
possible to do is to distinguish reasons for differences in the order of
magnitude of rate constants.
7.1
OBJECTIVES
The main aim of this unit is to study the Kinetics and mechanism
of redox reactions. After going through this unit you should be able to :

describe redox reactions and mechanism of one electron transfer
reactions;

explain outer sphere mechanism;

discuss theories of redox reactions;

describe potential energy diagram and explain Marcus theory;

discuss factors affecting electron transfer reactions; and

explain mechanism of inner sphere type reactions.
230
7.2
RE-DOX REACTIONS
The reaction in which the transfer of an electron from one atom to
the other occurs and hence the oxidation state of some atoms change is
called 'Redox reaction'.
These reactions of transition metal complexes are divided into two
classes on the basis of their mechanism :
1.
Electron – exchange Process : In which the electron transfer results
in no net chemical change. For example change of [Fe(CN)6]3-into
[Fe (CN)6]4- or that of [Co(en)3]3+ into [Co (en)3]2+.
These reactions have outer sphere electron-transfer mechanism and
are followed only indirectly, e.g. by isotopic lebelling or by nmr.
2.
Those reactions which involve net chemical change as a result of
electron transfer e.g. change of [Cr (NH3)5 X]2+ into [Cr (H2O)6]2+
or that of [Cr (H2O)5Cl]2+ into [Cr (H2O)6]2+.
These reactions follow inner sphere or bridge mechanism and can
be traced using standard chemical and physical methods.
7.2.1 Mechanism of one Electron Transfer Reactions :
Most of these reactions are believed to follow the following two
mechanism :
(a)
Electron transfer or outer – sphere mechanism, and
(b)
Bridge or inner – sphere mechanism.
7.2.2 Outer Sphere Mechanism
In these reactions neither the bonds are formed nor broken up.
Hence there is no change in the coordination sphere of metal ions, except
their oxidation states are changed.
231
These reactions occur by direct electron transfer and the electron
effectively hops from one species to the other and the ligands act as
electron-conduction media. It involves movement of an electron from the
outside of a ligand in one co-ordination sphere over the outside of a
second sphere. This mechanism is appropriate with large conjugated
ligands like phenonthroline and bipyridine.
For example : Transfer of an electron from [FeII(CN)4]4- to
[FeIII(CN)6]3-. Which can be studied by labelling of the complexes with a
radioactive isotope of Fe or 14C.
[*Fe2+(CN)6]4- + [Fe3+(CN)6]3-  [*Fe3+ (CN)6]3- + [Fe2+ (CN)6]4Ferrocyanide ion
Low spin and
inert
Ferricyanide ion
Low spin and
inert
Fe-C bonds Fe-C bonds
longer
shorter
In this reaction only charges of complex species are changed.
Fe(II) in ferrocyanide ion is oxidised to Fe(III), while Fe(III) in
ferricyanide ion is reduced to Fe(II) state. These reactions may be
considered analogues to collision – model.
The reaction is fast with second order constant 105 at 25oC and
there is no heat change in the reaction and gives same products after the
electron transfer both the axions are inert.
During this reaction none of the elements Fe, C, or N moves. The
Fe-C bond in [Fe (CN)6]3- is likely shorter than that in [Fe (CN)6]4-. Thus
if an electron is to be transferred between the axions in their ground state
equilibrium configurations by the Frank condon principle, the product [Fe
(CN)6]3- would be expanded and the [Fe (CN)6]4- would be compressed.
232
The reaction rates for outer sphere redox reactions, where the
complex retains its full coordination sphere, and the electron must pass
through both the coordination shells (a tunneling mechanism), vary from
(second order rate k2) 10-4 for [Co (C2O4)3]3- - [Co(C2O4)3]4- through 103
for MnO4- – MnO42- to more than 106 for [Fe (dipy)3]3+ - [Fe (dipy)3]2+
and [W(CN)8]3- - [W (CN)8]4- complexes.
Electron transfer processes where no change in the chemical
species takes place (e.g., the couples written above), can be followed by
the isotopic tracers or magnetic resonance methods. In one novel method,
the rate of loss of optical activity on mixing solutions of a D-complex of
one oxication state and the L-complex of another oxidation state (both
complexes being kinetically inert) gives the rate of electron transfer by
the reaction.
D–[Os (dipy)3]2++L–[Os (diply)3]3+ L-[Os(dipy)3]2+ +D–[Os (dipy)3]3+
Where both reactants are non-labile, e.g. in the case of [Fe (CN)6]4and [Fe(CN)6]3-, a close approach of the metal atoms is impossible, and
the electron transfer must take place by a tunnelling or outer sphere
mechanism. Although for an isotopic change the equilibrium constant is
nearly unity and Go is nearly zero, activation energy is required to
overcome the electrostatic repulsion between ions of like charge, to
distort the coordination of both species and to modify the solvent
structure around both species.
The [Fe (CN)6]4- - [Fe (CN)6]3- exchange reaction is catalysed by
alkali metal ions, the effect being greatest for caesium and smallest for
lithium. The very large cation Ph4As+ has little effect, however. These
results suggest that a partly desolvated cation accelerates exchange by
helping to overcome electrostatic repulsion by formation of a transition
233
state such as
[Fe (CN)6]4- ..... M+ ..... [Fe (CN)6]3A small cation holds its hydration sheath too strongly, whilst a very
large one does not bring the anions into close proximity. It is interesting
to note that the MnO42- - MnO4- exchange reaction is also subject to alkali
metal ion catalysis, the order of effectiveness being the same as for the
[Fe (CN)6]4- - [Fe (CN)6]3- reaction.
Outer-sphere reactions between complexes of different metals (e.g.
the [Os (dipy)3]2+ - [Mo(CN)8]3- - reaction mentioned earlier) are usually
faster than outer sphere exchange reactions between different oxidation
states of the same element. For such reactions the decrease in energy
when excited states of products are converted into ground states can
appear as the free energy of the reaction (Go must be negative, or the
reaction would not take place). This is tantamount to saying that for such
reactions the structure of the transition state is more like that of the
reactants; hence the activation energy is lowered and the rate is increased.
The outer sphere electron transfer may be represented as follows :
If oxidant = 0 and reductant = R, then
O+R

[ O .... R ]
[ O .... R ]

[ O .... R ]*
[ O .... R ]* 
[ O- .... R+ ]
[ O- .... R+ ] 
O - + R+
First the oxidising agent ( O ) and the reducing agent (R) come
closer to form precursor complex. During activation of the precursor
complex there is reorganisation of solvent molecules and the bond lengths
234
of metal ligand bonds are changed. However this takes place before the
electron transfer. In the last step ion – pairs are transformed in to product
ions.
For the outer sphere redox reactions, the following characteristics
are observed.
(1)
Electron transfer is expected to be fast when no change in
the molecular dimensions takes place.
(2)
Fast electron transfer takes place when the electrons are able
to reach the surface of the complex through conjugation (in
the ligands attached) or through monatomic ligands.
(3)
The ligand exchange is slower than the electron exchange
process.
(4)
The rate constant depends upon the cation present in the
solution; ion pair formation decreases the activation energy
by reducing the electrostatic repulsion energy.
(5)
7.4
Reactions involving large size differences proceed slowly.
THEORIES OF REDOX REACTIONS
Thus, for explanation of outer sphere reactions it is necessary to
consider too concepts : (i) Born – Oppenheir Approximation and (ii) the
energies of starting and end states.
According to Born–Oppenheir approximation, electron distribution
must be done, considering nuclei are stable at their place. If we consider
the nuclei are stable in the transition state, then we can see distribution of
wave function of the electron transferred on both the centres. On the basis
235
of energy, it will be more advantageous, the ion – ligand bond – length
stabilize at the intermediate value. At this point electron transfer takes
place at the bond – length of reactants.
According to the other concept, the possibility of electron transfer
will be maximum when the energy of the initial and final states are equal.
Consideration of these two concepts clearly show that the rate of
electron transfer and the activation energy for the process depends on the
capability of nuclei to reorganise.
An example of the reaction of inert reactants is electron transfer
reaction between solvated Fe (II) and Fe (III); which has been studied
using radio active isotope of iron (Fe*) :
[Fe (H2O)6]3+ + [Fe* (H2O)6]2+  [Fe (H2O)6]2++ [Fe* (H2O)6]3+
In this reaction also electron transfer takes place through outer
sphere or tunneling mechanism. At 25o second order rate constant is 3.0 L
mol-1 s-1 and the activation energy is 32 KJ mol-1. The arrangements
during electron transfer require changes in Gibbs energy; this is known as
inner sphere rearrangement energy *Gis. In addition to this, the energy
for the changes in the solvent outside the coordination sphere, i.e. Outer
sphere rearrangement energy, *Gos, is also important. Further, the
electrostatic interaction, energy between the two reactants, will be, *
GES. Hence the Gibbs energy of complete activation is the sum of all
these energies :
*G = *GIS + *GOS + *GES.
For electron transfer it is necessary that the energies of the
participating electronic orbitals are equal (Frank – Condon Principle). In
236
this reaction, one electron from t2g orbital of Fe (II) is transferred to t2g
orbital of Fe (III). The bond-lengths in Fe (II) and Fe (III) complexes are
different (In octahedral high spin complex of Fe (II) = 92 pm and that of
Fe (III) = 78.5 pm), which indicates the energies of orbitals are not
equivalent. If the electron transfer takes place without losing energy, then
we will get the product in which bond-length of Fe (II), complex is equal
to the characteristics bond – length of Fe (III) – complex and vice versa.
But this will be against the first law of thermodynamics. Actually, it is
necessary to give energy for electron – transfer. The actual reaction starts
with smaller bonds in Fe (II)-b. Complex and larger bonds in Fe (III) –
complex; unless the energy of participating orbitals become equal. The
vibrational stretching and compressions of metal – ligand bond help in
getting required configuration.
7.4
POTENTIAL ENERGY DIAGRAMS
Potential energy diagrams also confirm the relation between
molecular motion and electron transfer. For electron transfer it is
necessary that there should be coupling in vibrational and electronic
motion. The limit of these interactions is related with the energy
difference, E, at the crossing of potential energy diagrams (fig. 7.1). If
the coupling interaction is strong, distortion of bonds is very less and
electron transfer is easy. If the interaction is weak, bond-distortion in
high; G will be high and the reaction is slow. As the activated complex
rests at the intersection of two curves; while according to the noncrossing rule, molecular potential energy curves of same symmetry –
states do not cross each other, but divide into upper and lower curves
(Fig. 7.1). The indication of non-crossing rule is that, if the reactants
distort slowly in their ground states, then they are converted into
products, in their ground states itself, following the lowest energy path.
237
Fig. 7.1
In common red-ox reactions, Gibb's-energy of the reaction is not
zero. If the product surface is high (Fig. 7.2 (a)), then the crossing point
rises and the activation energy of the reaction will be high. On the other
hand, when crossing point drops (Fig. 7.2 (c)), activation energy
decreases. At the limit of exergonic reaction (Fig. 7.2 (d)) crossing point
rises up and the rate again slowdowns.
Fig. 7.2
In cyanide complexes of Fe (II) and Fe (III), according to LFT, the
extra electron of Fe (II) in t2g orbital, is in non-bonding molecular orbital;
and is diffused partially on the ligand due to  bonding. During electron
238
transfer between complex ions, [Fe (CN)6]4- and [Fe (CN)6]3- , initially
the products are in the exited state. If the shapes of reactant – species are
very less distorted, as compared to the transition state, then the energy of
activation required for electron exchange is very less and the reaction is
fast. Although, the value of equilibrium constant for isotopic exchange is
almost same and Go is zero; still energy of activation is required due to
following reasons :
(i)
For overcoming electronic repulsion between ions of same charge,
(ii)
to distort coordination shells of both the species, and
(iii)
to change the arrangement of solvent molecules surrounding the
two species.
7.5
MARCUS THEORY
Marcus proposed an approximate equation for rate constant of
outer sphere electron – transfer reaction. A perusal of fig. 7.2 clearly
indicates that the rate of electron transfer depends on two factors; one –
the shapes of potential energy diagrams and the other on the standard
Gibb's reaction energy. If the parabola in potential diagrams sharply rise,
with increase in bond strength, showing increase in energy, then crossing
points are high and also the energy of activation. On the other hand, less
deep potential curves show low activation energy. Similarly, bigger
values of equilibrium inter-nuclear distance show that the equilibrium
points are at longer distance. Hence, the crossing point will not be
obtained without big distortion. Higher is the value of standard reaction
Gibbs energy, lower will be the activation energy of the reaction. Based
on all these facts Marcus derived a relation for predicting rate-constant,
K, of outer sphere electron-transfer reaction :
K2 = f K1 K2 K
239
where, K1 and K2 are rate constants of the two exchange reactions
and K is equilibrium constant of the overall reaction.
f, is the function of rate constants and the interaction rate.
Generally, its value is taken 1 in all approximate calculations. As the log
of rate constants are proportional to activation energy, hence Marcus
equation may be expressed in the form of linear free energy relation :
2 In K = In k, + In K
i.e. 2* G1 + * G2 + t Go
Deviation from Marcus equation is supposed to be a special case in
outer – sphere reactions. These cases are due to the obstraction of high to
low spin during electron – transfer or due to change in the symmetry. An
important example, showing importance of bond distortion magnitude, is
self-exchange reaction of hexamine cobalt complexes :
[Co* (NH3)6]3+ + [Co (NH3)6]2+  [ Co* (NH3)6]2+ + [Co (NH3)6]3+
The rate constant of this slow second order reaction is 10-6 M-1 S-1.
The characteristics of the reactions are as follows :
(i)
Co – N bond Lengths, in Co (II) and Co (III) complexes, are quite
different, 2 11 pm and 194 pm respectively,
(ii)
Co (II) complexes are high spin complexes (t2g5 eg2) while Co (III)
complexes are low spin complexes (t2g6 ego); and
(iii)
after the electron – transfer, the configuration in both the
complexes probably becomes t2g2 eg1 and no ion remains in ground
state; i.e. they remain in excited state, resulting in increase in
activation energy. Because of high activation energy, this reaction
is quite slow, as compared to that between [Fe (CN)6]3- and [Fe
(CN)6]4240
Deviation from Marcus theory, otherwise considered as if the
reaction is not of outer sphere but is of inner sphere.
Generally, the outer sphere exchange reactions of different metal
ions in different oxidation states are faster, as compared to
reactions of same ions in corresponding oxidation states. An
example of the same is –
[Os (dipy)3]2+ + [Mo (CN)8]3-  [Os (dipy)3]3+ + [Mo (CN)8]4-.
The higher rate of these reactions is due to –
(i)
In these reactions when the excited state of the product
returns to the ground state, the decrease in energy, is
obtained in the form of free energy of the reaction (G
should be negative, otherwise reaction will not take place).
(ii)
It means, the structure of transition state is analogues to the
structure of the reactants, hence activation energy remains
low and rate increases.
Excited State Outer Sphere Electron. Transfer Reactions
Dramatic change are seen in the redox properties of
transition metal complexes, when they are in excited states,
absorbing energy. In this field, much work has been done with tris
(2-2 bipyridine) Ruthenium (II) Cation ([Ru (bpy)3)2+. Specially,
because this has possibility of decomposing water by photochemical reaction. Thus, this can open the path of the production of
hydrogen using solar energy. (Creutz and Sution, 1975).
hv
H2O
2+
[Ru* (bpy)3]
241
H2 + ½ O2
Here ruthenium complex function as photo-sensitizer. Go = 238
KJ mol-1
When [Ru (bpy)3]2+ absorbs 452 nm light, the very excited state of
initial [** Ru (bpy)3]2+ species changes into quite stable exited
species [Ru (bpy)3]2+ which in comparison to [Ru(bpy)3]2+,
(Ground State Species) is a better oxidant with 2.12 volt (+0.84 V
+ 1.28 V) and with -2.12 V (-0.86 V + 001.26 V) an outstanding
reducing agent (Fig. 7.3)
Fig. 7.3
The electronic transfer in the above absorption is a metal to ligand
charge transfer, in which one d-electron of ruthenium is exicited
and goes to  - antibonding orbital of a bipyridine molecule. Hence
in the excited state, the structure of [Ru* (bpy)3]+ may be written as
[Ru(III) (bpy)2 (bpy)]2+. The presence of an electron in an
antibonding orbital of a ligand makes this excited state cation quite
a better reducing agent, as compared the ground state cation.
Further the hole thus formed in ruthenium, increases its electron
accepting capability, resulting in the fact that this excited state.
cation, as compared to its ground state, is also a good oxiding
agent.
242
7.6
FACTORS
AFFECTING
REACTION RATE
ELECTRON
TRANSFER
Halpern summarised the factors affecting the rates of direct
electron – transfer reactions :
(i)
Electrostatic repulsion between ions of like charges increases the
activation energy; hence the rate of exchange of electrons
decreases.
(ii)
When there is no change in the shapes of the molecules, possibility
of fast electron transfer always remain.
(iii)
Fast electron transfer takes place, when due to coupling or by anatomic ligand, electrons reach on the surface of the complex.
(iv)
Generally, in comparison to ligand exchange, electron exchange is
fast. The value of rate constant depends upon cation present in the
solution. Due to formation of ion-pair, activation energy repulsion
decreases, because the electrostatic repulsion decreases, resulting
in increase in the rate.
(v)
With the increase in the conductivity of the ligand electron transfer
increases.
(vi)
When the difference in the shapes of oxidant and reductant are
much high, then there is possibility of slow reaction.
(vii) Higher is the negative value of Go for the reaction, faster is the
reaction.
243
Check Your Progress-1
Notes : (1) Write your answers in the space given blow.
(2) Compare your answers with those given at the end of the
unit.
a (i) Redox reactions involve ................................. from one atom
to the other.
(ii) Electron – exchange reaction involving no net chemical
change follow ................... electron transfer mechanism and
can be traced by ...................
(iii) The reaction in which electron – transfer results in net
chemical change follow ........................... mechanism and
can be traced by ..................
b (i) For electron – transfer it is necessary that energy of the
participating electronic orbitals ...................
(ii) During electron transfer the ....................... and ....................
of metal ligand bond help in getting required configuration.
(iii) The exited state electron transfer of ............................. opens
the path of the production of hydrogen using solar energy.
7.7
INNER SPHERE TYPE REACTIONS
Many oxidation-reduction reactions have been shown to occur by
a ligand-bridging or inner sphere mechanism in which substitution of the
coordination shell of one of the metal ions occurs. The classic example of
such a reaction is that between [Co (NH3)5 Cl]2+ and [Cr (H2O)6]2+ in
acidic solution, first investigated by Taube.
[Co (NH3)5 Cl|2+|Cr(H2O)6|2+ [(NH3)5 Co(III)–CI–Cr(II) (H2O)5.... (7.1)
[(NH3)5 Co – Cl – Cr (H2O)5]4+  [(NH3)5 Co]2++ [Cr(H2O)5 Cl]2+ .... (7.2)
[(NH)5 Co]2+ + H2O

244
[(NH3)5 Co(H2O)2+
..... (7.3)
[NH3)5 Co(H2O)2+ + 5H2O 
[Co(H2O]2+ + 5 NH3
...... (7.4)
The intermediate formed in reaction (7.1) dissociates to give a 6coordinated (Cr (III) and a 5-coordinated Co (II) Complex [reaction (7.2)]
which then picks up the additional water molecule from the medium
[reaction (7.3)] to develop into a 6-coordinated Co (II) complex. Being
unstable, the Co (II) complex undergoes complete equation to give the
hydrated Co (II) ion [reaction (7.4)].
This mechanism is supported by the following facts :
(i)
Chlorochromium (III) is formed during the reaction.
(ii)
No labeled * Cl atoms are found in the chlorochromium (III)
complex when the reaction is carried out in the presence of *Cl ions, indicating that no ionization of the complexes takes place.
If the reaction is carried out in the presence of
36
Cl- in the
solution, none of this isotope appears in the chromium (III) complex; this
fact provides further support for the bridging mechanism. Because the
change in oxidation states of the metal ions is accompanied by transfer of
a chlorine atom, the process is often referred to as an atom transfer
reaction.
Similar reactions occur when the chloride in the above example is
replaced by other halide ions, sulphate, phosphate, acetate, succinate,
oxalate and maleate. Among halide ions, the effectiveness for bridging
purposes (as measured by relative reaction rates) is F- < Cl- < Br- < I-, in
accordance with the expected order of ability to transmit an electron and
undergo covalent bond-breakage. For the organic ions mentioned, oxalate
and maleate (which contain conjugated systems) are considerably more
effective than acetate and succinate.
245
There are many other reactions which are believed to proceed by
the inner sphere mechanism. Where all the species involved are too labile
for tracer methods to be applicable or for the mechanism to be inferred
from the nature of the products, dependence of the rate of reaction upon
the concentration of an ion present in the solution can provide useful
information.
There is some evidence to show that change of anion makes
much more difference to the rates of inner sphere reactions (e.g. of [Cr
(NH3)5X]2+ - Cr2+ (aq), where X = F, Cl, Br or I) than to the rates of outer
sphere reactions (e.g. of the [Co (en)3]2+ [Co(en)3]3+ exchange catalysed
by F-,CI-or Br-or I-). It has therefore been inferred that the F-, CI- or Brcatalysed exchanges between Fe2+ and Fe3+, which proceed at about the
same rate, are all catalyse outer sphere reactions. This conclusion,
however, has been challenged in the case of the chloride ion catalysed
reaction, for which it is maintained that detailed interpretation of the
kinetics shows that the principal reaction involves atom transfer between
FaCl2+(aq) and Fe2+ (aq).
The consequences of the inner sphere mechanism of the redox
reactions are as follows :
(1)
Transfer of a ligand from one complex to the other.
(2)
The rate cannot be faster than the rate of exchange of the ligand
in the absence of the redox reaction.
(3)
The reaction is zero order with respect to one of the complexes
and of first order with respect to the other, where bond
dissociation takes place.
(4)
The reaction is first order with respect to the first species if the
rate determining step is the attack on the complex.
246
(5)
Electron transfer is rapid only if a conjugated bridged system is
formed in the intermediate.
(6)
Reactions involving large changes in the molecular sizes are
slow.
(7)
The rate of reaction between the Cr2+ and between CrX2+ and
between Cr2+ and [Co(NH3)5 X]2+ decreases in the order I- > Br>Cl> F- showing that the electron transfer through the bridged
halogen atom affects the reaction rate.
Check Your Progress – 2
Notes (1) Write your answers in the space given below.
(2) Compare your answers with those given at the end of the
unit.
(i) During inner sphere mechanism of redox reaction between
[Co(NH3)5Cl]2+ and [Cr (H2O)6]2+ the intermediate dissociate
to
give..................Cr
(III)
and
.................
Co
(II).
Complexes.
(ii) Amongst halide ions, the effectiveness for bridging purposes
is in the order ......<...........<.............
(iii) When the reaction is carried out in presence of *Cl- ion,
................ found in the chlorochromium (III) complex,
indicating that ................. of the complex takes place.
b (i) Inner sphere mechanism of redox reaction involve transfer of
................... to the other.
(ii) The reaction is ............... order with respect to one of the
complexes and of ................................... takes place.
(iii) Electron transfer is rapid only if .....................................
system is formed in the intermediate.
247
7.8
LET US SUM-UP
*
The reaction in which there is transfer of an electron from one
atom to the other occurs and hence the oxidation state of some
atoms change, is called 'Redox' reaction.
*
These reactions may be divided into two groups :
(a) The electron exchange process in which no net chemical
change takes place. These reactions have outer sphere
mechanism and are traced by isotopic labelling and nmr.
(b) The electron-exchange processes resulting in a net chemical
change are called inner sphere or bridge – mechanism redox
reactions. These may be traced by common standard chemical or
physical methods.
*
In outer sphere mecahnism reactions neither the bonds are formed
nor broken up, only the oxidation states of metal ions are
changed.
*
Where both reactants are non-labile eg. in [Fe(CN)6]4- and
[Fe(CN)6]3-, a close approach of the metal atom is impossible and
the electron transfer must take place by tunnelling or outer sphere
mechanism.
*
The [Fe (CN)6]4- - [Fe (CN)6]3- exchange reaction is catalysed by
alkali metal ions, the effect being greatest for caesium and
smallest for lithium.
*
Outer, sphere reactions between complexes of different metals
(eg. [Os (dpy)3]2+ - [Mo (CN)8]3- are usually fater than outer
sphere exchange reactions between different oxidation states of
the same element.
248
*
Electron transfer is expected to be fast when no change in the
molecular dimension takes place and the rate constant depends
upon the cation present in the solution; ion-pair formation
decreases the activation energy by reducing the electrostatic
repulsion energy.
*
According to the Born-Oppenheir approximation, electron
distribution must be done, considering nuclear are stable at their
place.
*
Accordingly, the possibility of electron transfer will be maximum
when the energy of initial and final states are equal.
*
For electron transfer it is necessary that the energies of the
participating electronic orbitals are equal (Frank – Condon
Principle).
*
Potential energy diagrams also confirm the relation between
molecular motion and electron transfer. For electron transfer it is
necessary that there should be coupling in vibrational and
electronic motion.
*
If the coupling interactions are strong, distortion of bonds is very
less and electron transfer is easy.
*
According to the non-crossing rule, molecular potential curves of
same symmetry-states do not cross each other, but divide into
upper and lower curves.
*
In common redox reactions, Gibb's energy of the reaction is not
zero. If the product surface is high, the crossing point rises and
the activation energy will be high.
249
*
Marcus proposed an approximation, giving an equation for rate
constant of outer sphere electron – transfer reaction.
Marcus relation is –
K2 = f K1 K2 K.
where K1 and K2 are rate constants of two exchange reactions and
K is equilibrium constant of the overall reaction; while f is the
function of rate constants and the interaction rate.
*
Redox properties of transition metal complexes in their excited
states are quite interesting. Important compound in this regard is
[Ru (bpy)3]2+, specially because this has a possibility of
decomposing water by a photo-chemical reaction.
*
Factors affecting electron transfer reactions are electrostatic
repulsion between ions of like charges, the shapes of the
molecules, electron reach on the surface of the complex, presence
of cation in solution, conductivity of the ligand and the values of
Go.
*
Many redox reactions occur by a ligand briding or inner sphere
mechanism e.g. [Co (NH3)5Cl]2+ - [Cr (H2O)6]2+ reaction. The
intermediate formed, dissociates to give a 6 – coordinated Cr (III)
and a 5 – coordinated Co (II) complex.
*
No labelled *Cl atoms are found in the chloro-chromium (III)
complex when the reaction is carried out in the presence * Cl
ions, indicating inner sphere mechanism and no ionisation of the
complex.
*
Amongst holide ions, the effectiveness for bridging purposes is in
the order F- <Cl- < Br- < I- in accordance with the e- transmitting
tendency.
250
*
The reaction is zero order with respect to one of the components
(complexes) and of first order with respect to the other, where
bond – dissociation takes place.
Electron – transfer is rapid only if a conjugate bridged system is
*
formed in the intermediate.
7.9
CHECK YOUR PROGRESS : THE KEY
1. (a) (i) transfer of electron
(ii) follow outer sphere
traced by isotopic labelling and nmr
(iii) follow inner sphere
traced by standard chemical and physical methods.
(b)
(i) are equal
(ii) the vibrational stretching and compression
(iii) [Ru (bpy)3]2+
2. (a) (i) Six coordinated Cr (III) and 5 – coordinated Co (II)
(ii) F- < Cl- < Br- < I(iii) no lablled *Cl atoms are
indicating that no ionisation
(b) (i) Ligand from one complex
(ii) Zero
and of first order
where bond dissociation
(iii) Conjugate bridge system.
251
Unit - 8
METAL-LIGAND EQUILIBRIA IN SOLUTION
Structure
8.0
Introduction.
8.1
Objectives.
8.2
Step-wise and Overall Formation Constants.
8.2.1
8.3
8.4
Thermodynamic Importance of Stability Constants.
Factors Affecting Stability.
8.3.1
Factors related with Metal.
8.3.2
Factors related with Ligands.
8.3.3
Chelate effect and its Thermodynamic Origin.
Methods of Determination of Stability Constants.
8.4.1
pH-metric method.
8.4.2
Spectrophotometric method.
8.5
Let Us Sum Up
8.6
Check Your Progress: The Key
252
8.0
INTRODUCTION
Generally complexes are designated as stable or unstable. The
general meaning of stability is supposed to be related with the concept,
whether a particular complex can be converted into other easily or not. As
a matter of fact, this is kinetic aspect of stability; which deals with the
rate of the reaction and its mechanism. The other aspect of stability is
thermodynamic aspect. In which stability of a complex is related with the
amount of energy released during its formation or the amount of energy
required to break it.
In this unit we describe complex forming equilibria in solution
and the various factors affecting it. We will also discuss the various
factors affecting stability constants for the formation of complexes in
solution. In the end of the unit we shall describe the method used for
determining stability constants of the complexes formed in solution.
Which involves quantitative characterisation of the complex-forming
reaction in solution.
You may recall what you have already studied about the basic
concept of chemical equilibria in solution.
8.1
OBJECTIVES
The main aim of this unit is to study the complex formation
equilibria in solution. After going through this unit you should be able to:
 describe stepwise and overall formation constants;
 explain thermodynamic importance of stability constants;
 discuss factors affecting stability of complexes; and
253
 describe methods of determining stability constants for binary
complexes in solution.
8.2
STEP-WISE AND OVERALL FORMATION CONSTANTS.
The term stability is a loose term, when the term stability is used
without qualification, it means that the complex exists and under suitable
conditions, it may be stored for a long time. The term can not be
generalised for complexes. A complex may be quite stable to one reagent
and may decompose readily in presence of another reagent.
In studying the formation of complexes in solution, two types of
stability of complexes is found:
1.
Thermodynamic Stability
This is a measure of the extent of which the complex will form or
will be transformed into another species under certain conditions,
when the system has reached in equilibrium. When we are
concerned with this type of stability, we deal with metal-ligand
bond energies, stability constant etc.
2.
Kinetic Stability
This refers to the speed with which transformation leading to the
attainment of equilibrium will occur. When we are interested in
kinetic stability for complex ions in solutions, we deal with rates
and mechanism of chemical reactions. These reactions may be
substitution, isomerisation, recemisation and electron or group
transfer reactions. In the kinetic sense, it is more proper to call the
complexes inert or labile complex rather than stable or unstable
complex. The complexes in which the ligands are rapidly
replaced by others are called labile, while those in which
substitution occurs slowly are called inert complexes.
254
Stepwise and Overall Formation Constants
According to J. Bjerrum (1941) the formation of a complex in
solution proceeds by the stepwise addition of the ligands to the metal
ion. Thus the formation of the complex MLn may be supposed to take
place by the following n consecutive steps.
where
M
=
central metal cation
L
=
monodentate ligand
n
=
maximum co-ordination number for the metal
ion M for the ligand
( ML)
M + L  ML K1 = [ M ][L]
ML  ML2 K2 =
ML2  ML3 K3 =
Thus
( ML2 )
[ ML][ L]
( ML3 )
[ ML2 ][ L]
MLn-1 + L  MLn Kn =
( MLn )
[ MLn1 ][ L]
The equilibrium constants, K1, K2, K3, ..........Kn are called stepwise
stability constants.
The formation of the complex MLn may also be expressed by the
following steps and equilibrium constants.
B
M + L 
ML,  =
1
( ML)
[ M ][ L]
B
M +2L 
ML2,  2 =
( ML2 )
[ M ][ L]2
B
Thus M + nL 
MLn,  n =
( MLn)
[ M ][ L]n
2
n
................(8.1)
The equilibrium constants,  1,  2,  3, ..........  n are called overall
formation or overall stability constants.  n is called as nth overall (or
cumulative) formation constant or overall stability constants.
255
The higher the value of stability constant for a complex ion, the
greater will be its stability. Alternatively 1/k values sometimes are
called instability constant.
Stepwise and cumulative stability constants are also expressed as
log10K1, log10K2................log10Kn and log10n respectively.
Relationship or Interaction Between n and K1, K2, K3, ..........Kn
K's and 's are related to one another consider for example, the
expression for  3 is:3=
( ML3 )
[ M ][ L]3
On multiplying both numerator and denominator by [ML] [ML 2]
and on rearranging we get:
3
=
=
=
Thus  n
=
=
or  n
[ ML3 ]
[ ML][ ML2 ]

3
[ ML][ ML2 ]
[ M ][ L]
[ ML ]
[ M ][ L]

[ ML3 ]
[ ML2 ]

[ ML][ L] [ ML2 ][ L]
K1 x K2 x K3
[ ML ]
[ M ][ L]

[ MLn ]
[ ML2 ]
.............
[ ML][ L]
[ MLn1 ][ L]
K1 x K2..........Kn
nn
=
K
n 1
n
From above relation, it is clear that the overall stability constant n
is equal to the product of the successive (i.e. stepwise) stability constants,
K1, K2, K3, ..........Kn. This in other words means that the value of stability
constants for a given complex is actually made up of a number of
stepwise stability constants.
256
8.2.1 Thermodynamic Importance of Stability Constants
In order to reach accurate conclusions regarding the nature of the
forces acting within complex species during their formation in solution,
the energy changes accompanying the reaction in question i.e. a complete
thermodynamic characterisation of the reactions is necessary at the very
least, determination of enthalpy ( H ), entropy ( S ) and free energy
( G )changes accompanying complexation.
In the language of thermodynamics, the equilibrium constant of the
reaction is a measure of the change in free energy, heat content and
entropy. A more useful manner of stating equilibrium constant is in terms
of the standard free energy change G , i.e. the difference of free energy
between the products and the reactants in a standard state, which is
related to equilibrium constants by the thermodynamic expression:
- RT log K = G = H - T S .....................................(8.2)
The reactions tends to go in the direction written, when G is
negative.
Enthalpy change ( H ) gives the amount of heat either consumed
or liberated per mole of products and is related to the strength of the
ligand to metal bonds, compared to that of the metal to solvent bonds.
Entropy change ( S ) is related to the change in randomness (the
disorder) of a system. As is quite evident from the relation given above
(8.2), complex formation is most favoured by the negative enthalpy and
positive entropy changes (either of the two or both) as may be expressed
by the equation:
log K =
S  H / T
......................................................(8.3)
2.303 R
257
In many reactions both the heat and entropy changes favour
complex formation but their relative importance changes markedly with
minor variations from ML to M'L or ML'.
8.3
FACTORS AFFECTING STABILITY
8.3.1 Factors related with Metal
The nature of the metal ions and the effect of the different physical
properties of the metal ions on the stability of the complex are:
1.
Stability (or stability constant) increases with decreasing size of
metal ion. K generally varies are 1/r.
2.
Stability constants for a complex increase with the charge of the
central ion. The K for the Fe(II) complexes will be less then the K
for the corresponding Fe(III) complexes.
3.
The ions with high polarizability give complexes with higher
stability constants. Thus Cu(I) complexes have higher K values
than the similar sized Na+ complexes, similarly of Ca2+ and
Cd(II) or Al (III) and Ga(III) the former have low K values for
the complex formation.
4.
Electronegativity increases the polarizing power and the ions with
higher electronegativity give stable complexes.
5.
Ionization Energies: The electronegativity, covalent nature and
ionic radii can be related to the ionization energies of the atoms.
It is found that the stability constants for the metal complexes
with a ligand increases with the ionization energies of the
metallic species.
Observations of Bjerrum Niecilson and others show that although
most of the metals of the periodic table form complexes, this tendency is
the most with transition metals. The reason being that the chelate effect is
258
almost an entropy effect for the metal ions of nontransitional group, while
for the transitions metals it is partly an enthalpy effect which increases
the crystal field strength. The increase in crystal field strength increases
the points of attachment of the ligand to the metal ion imparting greater
chelating tendency to the latter (cf. CFS). Fig 8.1
Fig. 8.1: CFSE affecting stability of aquo-complexes
Chatt Ahrland classified the metals into a and b classes while a
class metals form stable complexes with ligands having the coordinating
atoms, N, O, F (second period elements), b class metals form stable
complexes with ligands in which donor atom is P, S, Cl (third or latter
period elements).
The a class metals include H, alkali and alkaline earth metals; the
elements from Sc to Cr, Al to Cl, Zn to Br and lanthanides and actinides.
While amongst b class Rh, Pd, Ag, Ir, Pt, Au and Hg are included.
Elements from Mn to Cu, Tl to Po, Mo, Te, Ru, W, Re, Os, Cd are
border line metals.
It can be said with some approximation that increase in the ionic
charge of the metal ion and donor, will bring an increase in the chelating
tendency while the increase in ionic radius will decreases it. Thus small
259
cation size, comparatively large ionic charge and appropriate electronic
arrangements are responsible for the maximum ability of complex
formation by transition elements.
Mellor and Maley have shown that the stabilities of the complexes
of bivalent metal ions follow the order: Pd > Cu > Ni > Pb > Co > Zn >
Cd > Fe > Mn > Mg irrespective of the nature of the ligand. Irving and
Williams from the analysis of the data on stability constants of transition
metal ions, found that the order
Mn(II) < Fe(II) < Co(II) < Ni(II) < Cu(II) > Zn(II),
holds good. This order according to them follows logically from a
consideration of the reciprocal of ionic radius and second ionization
potential of the metal, and is known as 'Natural Order of Stability'.
Univalent ions have not been extensively studied but data on the
complexes of the univalent ions with dibenzol methanate ion shows the
order of the stability as:
Ag > Tl > Li > K > Rb > Cs
For tetravalent metals much less information is available, the
greater ease of hydrolysis of these ions making potentiometric titrations
more difficult. Irving and Williams suggest from a considerable limited
number of investigations that a rough order of stabilities be:
Ti > Fe > Ga > In > Al > Cr > Sc
8.3.2 Factors Related With Ligands
The properties of the ligands which affect the stability of the metal
complexes are as under:
260
1. Basicity of the ligands: The greater is the Lewis base strength,
higher is expected to be the stability constant of the complex. Thus
K values for the complexes are expected to change in a manner
similar to the changes in the proton association constant (BH) for
the ligands.
2. Dipole moment and polarizability of the ligands: Due to the
greater electrostatic interactions between the metal ion and the
ligands, polarity and ploarizability of ligand results in higher K for
the complexes.
3. (ML)  -bonding always increases the stability of the complex.
4. Steric factor: It play an important rule in determining the stability
constants for the complexes. Thus the 2 methyl derivative of 3
hydroxyquinoline gives much less stable complexes then the parent
compound because of the steric hindrance caused by the methyl
group adjacent to the site of co-ordination.
In complex formation hydrogen behaves just like a metal ion.
Therefore, a ligand with a larger affinity for proton will show the same
behaviour towards the metal ions. According to Riley any factor which
can increases the localization of negative charge in the co-ordinating
ligands makes the electron more readily available and thus increasing the
co-ordinating ability of a base. The correlation between the basic strength
of the ligand and the stability constant of the complexes was pointed our
first by Calvin and Wilson.
Ring Formation and Size of the Ring
Ring complexes or chelates are very stable due to reduced strain.
The number of ring formed, the size of the rings and stabilizing or
261
interfering resonance interactions are determined by the structure of the
chelating agent. The work of Ley on the chelates of amino-acids showed
that five and six membered rings are the most stable. Much evidence has
accumulated since then to prove that all chelates have either five or six
membered rings. Pfeiffer observed that in general the five membered
rings is the more stable when the ring is entirely saturated but when one
or more double bonds are present, the six membered rings is favoured.
Schwarzenbach and Co-workers have observed that there is a decrease in
clate stability with the increase in ring size. The stability of a five
membered ring is not chiefly due to entropy but rather to the enthalpy of
formation; the example being 1, 2, 3 triamine- propane tetra
chloroplatinum. Further the stability increases with the increase in the
number of rings in the molecule:
M(en) <
M(trien) < M(EDTA).
(one ring) (two rings) (five rings)
Steric Effect:
Steric hindrance can influence stability in many ways, e.g.
(i) Metal-ligand bonds are weakened due to the presence of bulky
group near the coordinating site.
(ii) The substituting group prevents the ligand from assuming the
planar configuration and hence introduce strain in the metal-donor
bond.
(iii)Steric hinderacne is also due to strained structure of the chelated
ring, since it breaks the usual linear configuration of the
complexes.
From the study of the copper complexes of substituted malonic
acids Riley concluded that ethyl and propyl groups had a larger effect
then methyl in reducing the stability.
262
Resonance Effects
The stability of a chelated ring will depend on the possibilities of
resonance in the ring and on how these will fit in with resonance in the
organic ligand itself. That resonance may affect the formation of a chelate
was first shown by Calvin and Wilson. The double bond resonance has
been attributed as a reason to be unusual stability of histamine cobalt
chelate.
Orbital hybridisation
There are certain factors which serves to make a specific bonding
arrangement stable. As an example, the shape of , ', ''
triaminotriethylamine is such that the bonding atoms must be grouped
tetrahedral round a metal atom. The ligand will therefore tend to form a
stable complex with a metal such a zinc, which favours sp 3 hybridisation
in its 4-co-ordinate compounds, rather than with one such as copper
which is limited to dsp2 (planar) hybridisation. Similarly, triethylene tetra
amine gives stable complex with metal ions having dsp 2 hybridisation,
rather then sp3 hybridisation.
8.3.3 Chelate Effect And Its Thermodynamic Origin
The chief factor responsible for the stability of the chelate ring is
the entropy change which can be viewed statistically or as probability
factor. Considering the electronic effect of the donor atom to be the same
in the monodentate and the bidentatc ligands, it can be seen that the
dissociation of a monodentate from a complex will be higher than that in
the chelating bidentate. The dissociation of the M-L bond in monodentate
will release the ligand completely from the coordination sphere of the
metal, so that it can be easily swept off by the solvent. But the
dissociation of one M-L bond for the bidenate ligand does not release the
ligand completely (for which simultaneous dissociation at both ends is
263
required). Hence the stability constant for metal chelate must be higher.
Consider the equilibrium reactions (Fig. 8.4):
[Co(NH3)6]3+ + 3en  [Co(en)3]3+ + 6NH3 ...................(8.4)
Assuming that (i) Co-N bond strength in the two complexes is
same (the f value of ammonia and ethylendiamine are within 3%), and (ii)
the entropy changes due to structure making and structure breaking are
negligible due to the similar size of the complexes, it can be seen that the
o
S will increase for the reaction as the number of moles of the products
are more than those for the reactants. This will help the reaction to go to
the right.
Check Your Progress-1
Notes : (i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
(i) Stability of metal complexes is primarily related with the
thermodynamic stability. Which deals with...........................
and....................
(ii) Overall stability constant, n for MLn complex is related with the
stepwise constants as- n = ..................................
(iii) The thermodynamic expression relating equilibrium constant is........................................
The reaction goes in the direction written when....................
(iv) CFSE results in the maximum increase in the stability of aquocomplexes of divalent metal ions in the first transition series at d n
configuration.......................and.........................
(v) The Irving Williams order of stability is..........................................
(vi) Chelate effect is primarily due to ............................ factor.
264
8.4
METHODS OF
CONSTANTS.
DETERMINATION
OF
STABILITY
There are many physical and chemical properties which may be
used to detect the formation of complex in solution and to measure the
stability constants. The detection of the complexes and the determination
of the stability constants are very closely related. Most of the methods
used for the detection of complexes can also be used to determine their
stability constants.
The study of the complexes is supposed to be incomplete without
finding the stability or formation constants, because most of the
properties and utility of the complexes depend on it.
The value of stability constants may predict the conditions required
for complete formation of a given complex. This knowledge of the
system is essential for correctly interpreting its optical and kinetic
properties of its partition equilibria and its biological behaviour.
Further, it may also help in planning analytical and separation
procedures. For example in case where the species is highly coloured or
can be precipitated from solution, extracted into an organic solvent or
absorbed on an ion exchange or chromatographic column.
Stability constant is related with the thermodynamic parameters, as
-RT, Ink, = G = H - T S
Where, G , H and S are changes of free energy of enthalpy
and of entropy respectively.
The stepwise or overall stability constant, thermodynamic
equilibrium constant gives the value of free energy change, associated
265
with the reaction. The corresponding changes on entropy change of
complex formation may be obtained by combining the stability constants
with the enthalpy change of complex formation, which is obtained by
determining the stability constant at a series of temperatures. The
knowledge of entropy is essential for the full understanding of many
factors such as size, shape, electronic structure of the central metal and
the ligand, the temperature and the composition of the solvent, which
influence the stability of the complex.
Let us consider a reaction between a metal M and ligand L to form
a complex MmLn.
Mm + nL  MmLn
K=
[ Mm Ln]
[ M ]m [ L ] n
where 'K' is stability constant of the complex MmLn. The stability
of the complex is quantitatively expressed in terms of dissociation
constant 1/k of the complex. The latter is the tendency of the complex to
split up into its components.
Some of the most important methods of determining the stability
constants are briefly described here.
8.4.1 pH - Metric Method
Bjerrum's Method
It is a potentiometric method for determining the stability constant
for complex formation. Although Bjerrum applied the method primarily
to the binding of simple molecules or negative ions to positive metal
ions. It may be used with equal success with chelating agents. The
theoretical relationship outlined by Bjerrum are not restricted to complex
formation but may be applied to any equilibrium process regardless of
266
the nature of the interacting substances. Thus, it has been used with
success on acid base, and redox equilibria. Although the reactions to be
considered involve ions that are more or less completely hydrated, rather
than the simple ions, but this fact does not affect the validity of the
conclusions, provided the activity of the water is maintained constant.
Formation or dissociation of a complex ion for molecule in the
solution always takes place in several steps, which can be easily
determined by measuring pH in this method.
Experimental Determination of Stability Constant by Bjerrum's
Method
This is a potentiometric method. When the lignad is a weak base or
acid, competition between hydrogen ion and metal ions for ligand can be
used to the determination of the formation constant.
Let us consider the equilibrium in which an acid and metal ions are
added to a basic ligand in solution. Thus the following equation are
obtained:
Ka
L + H+ 
HL+, Ka =
[HL ]
[L][H  ]
Basic Ligand Acid
KF
L + M+ 
ML+, KF =
[M L ]
[L][M  ]
Basic Ligand metal ion
Here Ka and KF are the acid association constant of the ligand and
formation constant respectively.
Now if CH, Cm and CL are the total amounts in moles/litre of acid
(H+) , metal (m+) and basic ligand (L), we have
CH = [H+] + [HL+]
CL = [L] + [ML+] + [HL+]
267
Cm = [M+] + [ML+]
Solving the last three equations given above and using the acid
association constant of the ligand, Ka. Then we get
[ML+] = CL - CH + [H+] -
C H  [H  ]
Ka [H  ]
[M+] = Cm - [ML+]
[L] =
C H  [H  ]
Ka [H  ]
Thus on putting the values of [ML+], [M+] and [L] from the above
equation in
K1 =
[M L ]
[M  ][L]
the value of K1 can be calculated. For the determination of [ML+],
[M+] and [L],
the values of CH, CL, Cm, Ka and [H+], is generally
determined potentiometrically using a PH meter.
In order to get better results, the ligand must be a medium weak
acid or base and the formation constant, K1, should be within a factor of
105 of the value of the acid association constant of the ligand, Ka.
Irving Rossotti Method
Calvin-Bjerrum pH titration technique as adopted by Irving &
Rossotti is generally used for determining the proton-ligand and metalligand formation constants. The procedure consists of:
(A) Determination of the formation curve of the system. This is

expressed as a plot of n (formation function) against pL for

metal ligand system and a plot of n A against pH for a proton-
268


ligand system (Definitions of the terms n , n A and pL are given
below).
(B) The calculation of the values of formation constants by solution
of the formation function of the system or otherwise.
(C) The
conversion
or
the
stoichiometric
constants
into
thermodynamic constants.

n term, was introduced by Bjerrum who called it the 'formation
functions' or 'ligand number ' and is defined as the average number of
ligand bound per metal atom or ion present in whatever form.

n = Total number of ligand ( L) bound to metal ( M )
Total number of M present in system
n
or

n
 . i [M Li]
=
i o
n
.....................................(8.5)
 .[M Li]
i o
which can be written using equation (8.1) as,
n

n
 . i .β[L]
=
i
[  = 1]................................(8.6)
i o
n
 .β
i o
1
[L]
i

A similar function for the proton-ligand sustems is n A, which
defined as the average number of protons bound per not complex bound
ligand molecule, and can be given by.
i

nA =
 . i .β
i o
i
 .β
i o
H
i
H
i
[H]i
[L]
[
H
o
= 1]........................(8.7)
i
whereas, pL gives the free ligand exponent and may be defined as.
pL = log .
1
[ L]
269
(A) Construction of the Formation Curves:
In Irving Rossotti method, this involves pH-titration of the
following three sets of mixtures (keeping total volume constant) against a
carbonate free standard alkali:
(a)
Mineral acid
(b)
Mineral acid + Ligand solution
(c)
Mineral acid + Ligand solution + Metal ion solution.
The ionic strength in each set is kept constant by adding
appropriate quantities of a neutral electrolyte solution. The temperature
of the solution in each case is kept constant. On plotting the observed pH
against the volume of alkali, one obtains (a) and acid titration curve, (b) a
ligand titration curve and (c) a metal-complex titration curve,
corresponding to the above titrations. [Fig. 8.2(a)]
The calculation of

n
are made from the volume of alkali required
to produce the same pH value in the metal and ligand titrations. Similarly

n
A values are calculated from the volume of alkali required to produce
the same pH value in the ligand and mineral acid titrations. According to


Irving and Ressotti, n A and n can be expressed as
nA =
Y TLo 
(V n  V 1 ) ( N  E  )
(V   V 1 )
TLo
iii 1
n
ο

n = (V  V ) (N  E )  TL o

(Vο  V 1 ) n
. TCM ο
....................(8.8)
....................(8.9)
Where Vo is the initial volume of the solution, Eo, TLo are the
initial concentrations of the mineral acid and the reagent respectively and
V', V'' and V''' are the volume of alkali of a given normality, N, required
270
during the acid, the ligand and the metal titration respectively at a given
pH (B). While the term Y gives the number of titrable hydrogen ions
arising from the chelating agent and TMo gives the initial concentrations
of the metal.
From the observed values of [L] for each

n
value, values of pL- are
calculated utilising the equation given by Irving and Rossotti:
n j
pL- log10

 nH (
n o
1
)n
anti log 

TCL0  n .TCM
0
.
V 0  V iii
Vo
....................(8.10)
Values of proton-ligand formation constants, K 1H , K 2H etc. obtained

from the proton-ligand formation curves plotted between values of n A
and pH [Fig. 8.2(b)].

The pH value at n A = 0.5 gives the value of log K 1H while the pH

value at n A = 1.5 gives the value of K H2 and so on.
Similarly, the values of stepwise stability constants of metalcomplexes are obtained from the formation curve plotted between the
values of n and pL- [Fig. 8.2(c)].
The value of formation constants are generally refined using least
square method.
Fig. 8.2: (a) pH - Titration Curves
271
(b) Proton-Ligand formation curve
(c) Metal-Ligand formation curve
8.4.2 Spectrophotrometric Method
1.
Job's Method
From the knowledge of stoichiometry of the complex, the value of
K (the stability constant) can be determined form the expression given
below, if the value of m and n are known:
K=
where, K
1/K
m n1  n m1  ( P  1) m n1 [n  (m  n) x]
m  n 1
C1
 P n1 [ P(m  n) x 1 ]mn
=
Stability Constant
=
Dissociation Constant of the complex.
272
P
=
Ratio of the concentration of the ligand to the
concentration of metal.
C1
=
Molar concentration of metal solution.
X
=
Concentration of ligand for which the concentration
of complex is maximum.
m
=
The number of moles of a metal required to combine
with "n" moles of ligand.
for (1:1) Metal ligand ratio in the complex
m=n=1
K=
( P  1)(1  2 x)
C1  [ ( P  1)( x  1)]2
Vosburgh and Cooper as well as Katzin and Gebert have extended
Job's treatment to systems in which two or more complexes are formed.
The ratio of the concentration of metal should not be equal to 1 i.e. nonequimolecular solutions of ligand and metal should be used.
2.
Turner Anderson Method
Turner and Anderson have modified Job's method and have
successfully used for determination of stability constant. By plotting a
continuous variation curve for a given range of compositions and then
repeating the procedure for more dilute solutions. If the initial
concentrations of the metallic ions and ligands are 'a' and 'b' respectively,
then
K=
X
(a  x)(b  x)
where, K
=
Stability Constant
X
=
Concentration of the complex
273
It is assumed that Beer's Law is obeyed, i.e. the optical density of
the solution is proportional to the concentration of the complex in the
given range. If, therefore, any two solutions on the two curves have the
same optical density, as shown in the graph a1, a2 and b1, b2 represent the
concentrations of the metal and the ligand respectively on the two curves,
then:
K=
X
X
=
(a1  x)(b1  x) (a2  x)(b2  x)
Where, the subscripts 1 and 2 refer to the reagent concentrations.
Thus K be calculated by solving the equation.
Molar Concentration
Fig. 8.3 :
Deskin has extended the method to the study of complexes formed
in the ration of 1:2, then:
M + 2L = ML2
K=
X
(a  x)(b  2 x) 2
Taking the concentration a1, a2 and b1, b2 for the same absorbance
i.e., the same value of x, we have
K=
X
X
=
2
2
(a1  x)(b1  2 x)
(a2  x)(b2  x)
The value of x is determined from the relation
4x2(a1-a2+b1-b2)-x(4a1b1-4a2b2+b 12 - b 22 )+ (a1b12-a2b22) = 0
274
i.e. AX2 + BX + C = 0
Where, A
=
-B =
4(a1 -a2 + b1 - b2)
4(a1b1 - 4a2b2 + b 12 - b 22 )
or b1(4a1 + b1) - b2(4a2 + b2)
C
=
(a1b 12 - a2b 22 )
By solving the quadratic equation:
( B)  ( B 2  4 AC )
K=
2A
( B)  ( B 2  4 AC )
K=
2A
By knowing the value of X, the value of K can be calculated.
Similarly, if metal and ligand react in the ratio 2:1 then.
2 M + L = M2 L
Taking the concentration a1, a2 and b1, b2 for the same absorbance
i.e., for the same value of X, we have
K=
X
X
=
2
(a1  2 x) (b1  x) (a2  2 x) 2 (b2  2)
or 42(a1-a2+b1-b2)-(4a1b1-4a2b2+a 12 - a 22 + (a 12 b- a 22 b2) = 0
i.e. AX2 + BX + C = 0
Where, A
=
4(a1 -a2 + b1 - b2)
-B
=
(4(a1b1 - 4a2b2 + a 12 - a 22 )
or b1(4a1 + b1) - b2(4a2 + b2)
C
=
(a 1 b 1 - a 12 b 2 )
By solving the quadratic equation, the value of X is determined
K=
or X =
( B)  ( B 2  4 AC )
2A
( B)  ( B 2  4 AC )
2A
Mushran has modified this method so as to suit for 1:3 complexes.
275
3.
Mole Ratio Method
The you and jones method can also be utilised for determination of
the stability constants.
Fig. 8.4
The extrapolated value (A extp.) (fig. 8.4) near the "equivalence
point" on the plots correspond to the total absorbance of the complex. If
the complex formed is complete. Actually the complex is slightly
dissociated in this region, and the absorbance read is somewhat low. The
ratio of the true absorbance to the extrapolated absorbance is the mole
fraction of the complex actually formed.
A
[mx]

A extp
c

where
c
is the total analytical concentration (expressed in moles/litre) of

the metal or ligand, whichever has the limiting concentration at the point
in question. Therefore
[MX] = A/A extp. C
M = Cm - (mx) = Cm (A/A extp. ) C
X = Cm - (mx) = Cx - (A/A extp.) C
276
Stability constant K =
 K=
[ MX ]
[ M ][ X ]
[(A / A extp. ) C]
[(Cm  A / A extp. ) C ] [(Cx  A / A extp. )C ]
-
Where A
=
Absorbance at the metal ligand ratio.
A extp.
=
The extrapolated value of Absorbance.
Cm
=
Concentration of the metal at equivalence point.
Cx
=
Concentration of the ligand at equivalence
point.
C
=
Total analytical concentration of the ligand.
When metal ligand ratio and the ratio shown by extrapolation do
not be on the same ordinate, then the value Cx and C will not be the
same. C is calculated at the point of intersection of the extrapolated
curve.
4.
Raghav Rao's Method
Subbarama Rao and Raghav Rao used job's method of continuous
variation and molar ratio method for determination of stability constants.
They used equimolar solutions of metal and ligand with optical density as
he index property. This method is also known as graphical method.
Reddy and Seshoish used the same graphical method using
conductance and optical density as the index property.
277
Check Your Progress-2
Notes : (i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
(i) Irving-Rossotti method is a modification of.......................method.

(ii) n is called.......................................and is defined as ...................
........................................................................................................
(iii) pL- = .........................................
(iv) Formation-curve is a plot between........................and.....................
(v) Turner Anderson method is a modification of ........................
method used for determination of ................................. by
plotting ............................... curve for a given range of .................
(vi) The extra plotted value in the mole ratio plot near the equivalence
point corresponds to ............................................. the complex.
8.5

LET US SUM UP
Stability of complexes in aqueous solutions is related with the
thermodynamic aspect, which deals with metal-ligand bond energy
and stability constants.

The formation of MLn complex in solution is supposed to take
place in n steps. In each step one mole of ligand is bound with the
metal ion replacing a mole of the coordinated water.

The equilibrium constants K1, K2, K3, ..........................Kn for the
reaction in each step of the complex formation are known as
'stepwise formation constants' and are related with the 'overall
stability or formation constant' n , i.e. the equilibrium constant for
the overall reaction: M + nL  MLn,
278
as : n = K1, K2, K3, ..........................Kn
nn
or n   K n
n 1

The equilibrium constant is related to the thermodynamic
expression as follows:
- RT log K =  G =  H - T  S .

The factors affecting stability of complexes are mainly related with
the metal ion and the ligands.

The factors due to metal are primarily related with the size of the
ion, its charge, possibility of  -bonding and CFSE gained.

Stability is proportional with the charge and ionic potential (e/rratio) but is inversely proportional with the size of the metal ion.

ML,  -bonding increases it, while LM,  -bonding decreases
it.

Similarly higher is the CFSE higher will be the stability.

a-groups metal form stable complexes with ligands N, O, F doner
atoms; while b groups metals give more stable complexes with the
ligands, having P, S and Cl donor atoms.

The Irving Williams order of stability is:
Mn (II) < Fe (II) < Co(II) < Ni(II) < Cu(II) > Zn(II).

Factors related with the ligands are mainly basicity, dipole moment
and polarizability of ligands, possibility of  -bonding and steric
factor.

Stability is proportional with the basicity, dipole and polarizability
of ligands.

ML,  -bonding (complexes with the unsaturated ligands)
increases the stability.
279

Shape of the ligand molecule also affects stability e.g. while
triethylene teramine gives complex with metal ions having dsp2
hybridisation(sq.
planar
geometry);
,
I
 ,

II
triamminotriethylamine gives stable complex with metal ions
having sp3 hybridisation (Tetrahedral geometry).

Chelates are more stable compared to non chelates.

Stability increases with the number of rings formed per mole of the
ligand e.g.
M(en) < M(trien) < M(EDTA)
(1 ring)

(2 ring)
(5 ring)
Higher stability of chelates is mainly related with the entropy
factor.

The stability constants of metal complexes in solution are
determined generally using two methods: one the potentiometer
(pH) titration method due to Bjerrum and its modification by Irving
and Rossotti; and the other one spectrophotometer methods due to
job and its modification by Turner-Anderson.

In Irving Rossotti method stability constants are computed by

plotting formation curves, between n (the formation function) and
pL-.


n is the average number of ligand bound per metal atom or ion;
while pL is the free ligand exponent; log

I
[ L ]
According to half integral method:

The value of pL- at 0.5 n = Log K1

The value of pL- at 1.5 n = Log K2 and so on.

The values of stability constants are generally refined by least
square method.
280

Turner and Anderson method involves plotting a continuous
variation curve for a given composition and repeating the
procedure for more dilute solutions.
8.6
CHECK YOUR PROGRESS: THE KEY
1
(i)
Deals with M-L bond energy and stability constants.
(ii)
Related as βn   K n
n n
n 1
(iii) - RT log K =  G =  H - T  S .
(iv)
 G is negative.
(v)
Mn (II) < Fe (II) < Co(II) < Ni(II) < Cu(II) > Zn(II)
(vi) Entropy factor.
2
(i).
Bjerrums's method
(ii)
Formation function, defined as the average number of ligand
bound per metal or ion.
I
[ L ]
(iii)
PL- =
(iv)
Between n and PL-
(v)
Job's method used for determination of stability constants by

plotting continuous variation curve.............of composition.
(vi)
The total absorbance of.
281
Unit - 9
METAL CLUSTERS
Structure
9.0
Introduction.
9.1
Objectives.
9.2
Boranes and Higher Boranes.
9.3
9.4
9.2.1
Wade's Rule.
9.2.2
Closo-Boranes.
9.2.3
Nido-Boranes.
9.2.4
Arachno-Boranes.
9.2.5
Structural Interrelation.
9.2.6
Synthesis.
9.2.7
Reactions.
Carboranes.
9.3.1
Synthesis.
9.3.2
Properties.
9.3.3
Structures.
Metalloboranes and Metallocarboranes.
9.4.1
Properties.
9.5
Metal Carbonyl Halides.
9.6
Compounds with metal-metal multiple bonds.
9.7
Let Us Sum Up.
9.8
Check Your Progress: The Key.
282
9.0
INTRODUCTION
Closed polyhedrons play important part in the synthesis of cluster-
molecules in inorganic chemistry. These cluster-molecules include
polyhedral boranes, carboranes and metalloboranes and metallo
carboranes; organometallic clusters and metal halide clusters.
The definition of metal clusters includes those molecular
complexes in which metal-metal bonds form a triangular or a large closed
structure. This definition does not include linear M-M-M bonded
compounds or those cage like structures in which metal atoms, in closed
structures are interlinked through ligands, forming M-L-M bonds.
Presence of metal-metal (M-M) bond in these molecules may be
ascertained with the help of data of bond lengths and also the stability of
compounds. As amongst d-block groups, metal-metal bond strength
gradually increases moving down a group, hence d-block metal in fourth
and fifth periods of the periodic table form M-M bonded compounds in
large number.
9.1
OBJECTIVES
The main aim of this unit is to study the nature, methods of
preparation and structures of metal-clusters. After going through this unit
you should be able to:
 describe boranes and higher boranes with reference to their
classification, synthesis reactions and structures;
 discuss carboranes and explain their synthesis and properties in
the light of their structures;
 describe metalloboranes and metallocarboranes in relation with
carboranes;
 explain structures of metal carbonyl halides; and
283
 identify compounds with metal-metal multiple bonds and their
structures.
9.2
BORANES & HIGHER BORANES.
Boron hydrides are known as Boranes. These are named boranes
in analogy with alkanes. These are gaseous substance at ordinary
temperatures.
It is expected that boron would form the hydride BH3, but this
compound is unstable at the room temperature. However, higher hydrides
like
B2H6(diborance).
B4H12
(tetraborane),
B6H10(hexaborane),
B10H14(decaborane) etc. are known. The general formula of boranes are
BnHn + 4 and BnHn + 6 (Proposed by stock). In addition to these is one,
recently discovered series of closed polyhedral structures with the
formula [BnHn]2-. Higher boranes have different shapes, some resemble
with nests, some with butterfly and some with spider's web.
The modern explanation of the structure of boranes is due to
C.L.Higgins, who proposed the concept of three centred two electron
bond (  -bond) Fig. 9.1. He also proposed the concept of completely
delocalised molecular orbitals to explain structures of boron polyhedrons.
He established icosahedral structure of [B12H12] Fig 9.2.
Fig. 9.1: 3C, 2e bond in B2H6
284
Fig. 9.2: B12H12 Icosahedron
In higher boranes, in addition to two centred two electron (2c, 2e)
and
the three centred two electron bond (3c, 2e bond) present in
diborance, B-B 2C, 2e and B-B-B (3c, 2e) bonds are also important. In BB-B bonds, three atoms of boron with their sp3 hybridisation are placed at
the corners of a equilateral triangle (Fig. 9.3).
Fig. 9.3: B-B-B bond
9.2.1 Wade's Rule
In 1970 K. Wade gave a rule relating the number of electrons in the
higher borane molecules with their formulae and shapes. Using these
rules one can predict the general shapes of the molecules from their
formulae. These rules are also applicable on carboranes and other
polyhderal molecules called 'Deltahedral's Deltahedrons are so called, as
they are composed of delta,  , shaped triangular faces.
285
According to Wade's rule, the building blocks of deltahedrons are
BH units, which are formed by sp-hybridisation of boron atom. Out of the
two sp hybrids one is used in the formation of 2c, 2e B-H exo bond of the
deltahedron and the other sp hybrid is directed inside as a radial orbital.
Remaining two unhybridised p orbitals of each boron atoms are placed
perpendicular to the radial orbitals and are known as tangential orbitals.
These radial and tangential orbitals combine by linear combination
method to form skeleton or framework of the deltahedron. To fill all
bonding molecular orbitals of the skeleton, necessary number of electrons
are obtained form the radial orbitals of BH units and s orbitals of the extra
hydrogen atoms. These electrons are called Skeletal electrons. For
example in B4H10, four BH units contribute 8 electrons (4x2 = 8) and six
extra hydrogens give six electrons thus B4H10 has total 14 skeletal
electrons Fig 9.4 gives the molecular energy diagram of [B 6H6]2-. This
molecule has seven pairs of skeletal electrons (six boron atoms and one
pair from two negative charges). These are used to saturate seven skeletal
molecular orbitals (a1g, t1u and t2g).
tui
eg
t2g
t2u
t2g
t1u
a1g
Fig. 9.4: Skeletal molecular energy diagram of [B6H6]2-
286
Classification:
On the basis of structures, molecular formula and skeletal electrons
higher boranes are classified into Closo, Nido, Arachno and Hypo (Table
9.1):
Table 9.1
Closo
[BnHn]2-
Skeletal
Electron Pair
n+1
Nido
[BnHn+4]
n+2
B2H6 , B5H9, B6H19
Arachno
[BnHn+6]
n+3
B4H10 , B5H11
Hypo
[BnHn+8]
n+4
Only derivatives are known.
Name
Formula
Examples
[B5H5]2- to [B12H12]2-
9.2.2 Closo Boranes
These are closed structured (Closo, Greak, meaning cage) boranes
with the molecular formula [BnHn]2- and skeletal electrons = n+1 pairs (=
2n+2 electrons). In this structure, there is one boron atom placed at each
apex and there are no B-H-B bonds present in the molecule. All the
member of the series from n=5 to 12 are known. [B 5H5]2- is trigonal
bipyramidal, [B6H6]2- is octahedral and [B12H12]2- is icosahedral. All are
stable on heating and are quite inert.
9.2.3 Nido-Boranes
These boranes have nest (Nido, Latin, meaning Nest) like structure.
Their general formula is BnHn+4 and have (n+2) pairs = 2n+4 skeletal
electrons on removing one boron atom from an apex of closo structure,
nido structure is obtained. Because, of the lost boron atom, these boranes
have extra hydrogens for completing the valency. The polyhedra in this
series have B-H-B bridge bonds in addition to B-B bonds. They are
287
comparatively less stable than 'Closo', but more than 'Arachno' on
heating.
9.2.4 Arachno-Boranes
These boranes have the general formula (BnHn+6) and skeletal
electrons = (n+3) pairs = 2n+6 = electrons. These molecules are obtained
by removing two boron atoms from two apexes of the closo structure and
have spider-web like structure. They have B-H-B bridge-bonds in their
structures and are very reactive and unstable on heating.
9.2.5 Structural Inter-relation
The structural interrelation between closo, nido arachno species is
shown in Fig. 9.5.
Arachno B4H10
Fig. 9.5:
288
This is based on the observation that the structures having same
number of skeletal electrons are related with one another by the removal
of BH unit one by one and the addition of suitable number of electrons
and hydrogen atoms, e.g. by removing one BH unit and two electrons
from octahedral closo. [B6H6]2- ion and adding four hydrogens, we get
square pyramidal nido- B5H9 borane. On repeating same process on nido
B5H9 (i.e. removing one BH unit and adding two hydrogen's), we get
butterfly shaped arachno. B4H10. Each of these three boranes have 14
skeletal electrons, but due to removal of BH unit, the resulting structure
becomes more open gradually (Fig. 9.5). The most symmetrical closo
structure has (n+1) skeletal molecular orbital, which requrie 2n+2
electrons. Similarly, nido-boranes have (n+2) molecular-orbitals and need
2n+4 skeletal electrons; while for (n+3) molecular orbital, arachno
boranes require 2n+6 skeletal electrons (see fig 5.6 for comparison
between these classes of boranes).
9.2.6 Synthesis
The simplest method for synthesis of higher boranes is the
controlled pyrolysis of diborance, B2H6 it is a gas phase reaction, BH3
formed in the first step reacts with borane to give higher boranes:
B2H6(g)

2BH3(g)
B2H6(g) + BH3(g)

B3H7(g) + H2(g)
B3H7(g) + BH3(g)

B4H10(g)
B2H6(g) + BH3(g)

B3H9(g)  [B3H8]-(g) + H+
5[B3H6](g)

[B12H12]2-(g) + 3[BH4]-(g) + 8H2(g)
2[BH4](g) + 5B2H6(g)

[B12H12]2- + 13H2
289
Closp
Nido
Arachno
Fig. 9.6: Interrelation between closo, nido and arachno-boranes
9.2.7 Reactions
The important reactions of higher boranes are with Lewis bases,
which involve removal of BH2 or BHn from the cluster, growth of the
cluster or removal of one or more number of protons:
1.
Decomposition by Lewis-bases:
B4H10 + 2NH3  [BH2(NH3)2] + [B3H8]
The reaction is analogous to the reaction of diborane with
ammonia.
290
2.
Deprotonation :
Higher boranes give deprotonation reaction easily rather than
decomposition:
B4H10 + N(CH3)3 [HN(NH3)3] + [B10H13] This deprotonation takes place from 3c, 2e BHB-bond. The
bronsted acidity of boranes increases with their size:
B4H10 < B5H9 < B10H14
For deprotonation of B5H9 strong-base like Li4(CH3)4 is
required:
B5H9 + Li(CH3)  Li+[B5H8]- + CH4
3.
Cluster Building:
Reactions of borane with borohydride are important with
respect to synthesis of higher boranes:
5K[B9H14] + 2 B5H9  5K[B11H14] + 9 H2
4.
Electrophilic displacement of proton:
Electrophilic displacement of proton by the catalytic activity
of Lewis acids like AlCl3 is the basis of alkylation and
halogenation of boranes:
AlCl
 [CH3B5H8] + HCl
B5H9 + CH3Cl 
3
291
Check Your Progress-1
Notes : (i) Write your answers in the space given below .
(ii) Compare your answers with those given at the end of the
unit.
A(i) Metal Clusters include those molecular complexes in which
..............bonds form a....................or large ...............................
(ii) Higher boranes may have different shapes resembling (a)
(b)
(c)
(iii) The various types of bonds present in higher boranes are mainly(a)
(b)
B(i) Wade's rule relates (a)
(b)
and (c)
(ii) Main polyhedral structure of higher boranes is called
............................. which have ....................... units as the building
blocks.
(iii) Main classes of higher boranes with their general formula and
skeletal electrons pairs are Name
Formula
Skeletal electron pairs
(a) ......................... ...............................
.............................
(b) ......................... ...............................
.............................
(c) ......................... ...............................
.............................
292
9.3
CARBORANES
Carboranes are mixed hydrides of carbon and boron, having both
carbon and boron atoms in an electron - deficient; skeletal framework.
There are two types of carboranes:
1. Closo-Carboranes: These have closed cage structrues in which
hydrogen bridges are structurally analogous to the Bn Hn-2 anions
with B- replaced by isoelectronic carbon. These carboranes have
the general formula. C2Bn+2 (n=3) to 12. The important member is
C2B10H12 (Fig. 9.7). Which is isoelectronic with [B12H12]2- similarly
B4C2H6 is isoelectronic with [B6H6]2-.
(A) 31, 2, C2 B10 H12
(B) C2B4H6
Fig. 9.7
2. Nido Carboranes: They are having an open case structure in
which some framework members are attached likely by hydrogen
bridges. These are derived formally from one or other of several
borones.
These contain one to four carbon atoms in the skeleton.
In addition to the above types of carboranes, there are a
number of carboranes with an additional heteroatom such as
phosphorus built into the basic structure and a family of metallo
carboranes, some of which are similar to ferrocene. One peculiar
feature common to all carboranes is that to date no compound has
293
been synthesized with either carbon bridging two boron atoms in a
three centre two electron bond or acting as one end off a hydride
bridge.
First carborane was obtained in 1953 when mixtures of
diborane and acetylene were ignited with a hot wire. Since that
time, many new carboranes have been isolated.
Nomenclature:
Rules for naming carboranes are as follows:
i. First of all, give the positions and number of carbon atoms, then
the type of carborane (either closo or nido) and finally the name of
the borane from which the carborane is formally derived and the
number of hydrogen atoms shown in bracket. For example CB5H9
is name as monocarbonido hexaborane (9). Similarly, the three
isomers of C2B10H12 are named as 1, 2; 1, 7 and 1, 12 dicarbocloso-dodecaborane (12).
ii. Number of atoms in these structure are counted by starting the
numbering from that in the apical position and proceeding through
successive rings in a clockwise direction.
This rule is important in naming the isomers.
Closo-Carboranes or Closed Cage Carboranes
These carboranes are having general formula C2BnHn+2 (n=3 to 10)
in which the constituents are only terminal. These are isoelectronic
with the corresponding [BnHn]2- ions and have the same closed
polyhedral structures, with one hydrogen atom bonded to each
carbon and boron. No bridging hydrogen atoms are present in the
C2Bn skeleton. They are considered in three groups.
a.
small, n = 3 - 5
b.
large, n = 6-10 and
294
c.
dicarbo-closo-dodecaborone
9.3.2 Preparation:
I(a)
The Small Closo Carboranes (C2BnHn+2 where n = 3 to 5)
C
B5H9 + C2H2 490

 1,5 - C2B3H5 + 1,6 - C2B4H6 +2,4 - C2B5H7
o
Example - The closo hexaborane isomers, C2BnH6,
(b)
The Large Closo Carboranes (C2B2Hn+2 where n = 6 to 9)
The first three members of this group of carboranes are
obtained by the thermolysis of 1,3 - C2B7H13 and 1,3 - C2B2H12.
Example : C2B6H8 is made from hexaborane (10) and
dimethylacetylene. The structure of 1,7 - Me2C2B6H6 is based on
the bicapped triangular prism. The carbon atoms are present one on
the prism and the other above the face opposite.
(c)
Dicarobo-closo-dodecaborone:
Preparation: The orthocarborane is the only isomer which
can be synthesized directly. However, it is synthesized by the base
catalysed reaction of acetylenes with decarborane (14) or via
B10H12L2.
H
RC
B10H14 + 2L 

 B10H12L2 
 R2L2B10H10 + H2 + 2L
2
2 2
Example: C2B10H12 gives three isomeric structure - 1,2 (ortho), 1-7
(meta) and 1, 12 (para)
(II) Nido-Carboranes or Open Cage Carboranes
These structures are derived formally from one or other of
several boranes and contain from one to four carbon atoms in the
skeleton.
Examples: CB5H9, C2B4H8, C3B3H7, C4B2H6 etc.
Preparation: The smaller nido-carboranes are generally prepared by
reacting a borane with acetylene under mild conditions.
295
Example: B5H9and C2H2 undergo reaction in the gas phase at 215oC
to give mainly the nidocarborane 2,3 - C2B4H8 together with methyl
derivatives of CB5H9.
The preparation method described above does not yield a
single product but a mixture of several products whose separation is
not an easy task. However some smaller nidocaroranes are prepared
by the following specific methods:
i.
Mono carbo-nido-hexaborane (7) CB5H7 is formed by
passing
silent
electric
discharge
through
1-methyl
pentaborane (9).
ii.
The only example isoelectronic with B5H9 is 1,2dicarbonido - pentaborane(7), C2B3H7, which is prepared as
follows:
C


 C2B3H7 (3 - 4 % yield)
B4H10 + C2H2 50
o
iii. Monocarbonidohexaborane (9), CB5H9 is formed from
ethyldifluoroborane and lithium.
The nido-carboranes are formally related to B6H10. All are having
eight pairs of electrons which are bonding the six cage atoms together.
Large Nido-Carborane:
Dicarbo-nido-undecaborane, C2B9H13, is the second member
of the class of nido-carboranes C2BnHn+4 (n =4 or 9),. The parent
carborane and its substituted derivatives can be prepared by the
base
degradation
of
ortho-carborane
(1,2-dicarbocloso-
dodecaborane (C2B10H12).

1
MeO
H
 C2B9H12 
 C2B9H13
1,2 - C2H10H12 
When C2B9H13 is heated, the closo-undeca-Borone (11) cage
is formed.
296
9.3.2 Properties
Properties of carboranes resemble with that of the corresponding
boranes closely. Thus, 1.2 dicarbo closo-dodecarborane-12 is stable in
both air and heat. On heating in inert atmosphere at 500 oC, it is converted
into 1, 7 isomer i.e. meta or neo isomer; while at 700 oC it is concerted to
1, 12 isomer i.e. para-isomer (Fig. 9.8)
Fig. 9.5: (a) C2B10H12
(b) 1,7 C2B10H12
(c) 1,12 C2B10H12
Analogous to boranes, carboranes are also classified into closo,
nido and arachno structure.
The chemical reactions, in so far as they are known, are very
similar to those of C2B10H12, which are described below. Various
substitution reactions have been studied and the hydrogen atoms bonded
to carbon are weakly acidic.
All three of the icosahedral isomers are stable both to heat and to
chemical attack, and much more so than decaborane (14). They are white
crystalline solids which resist both strong oxidizing agents and strong
reducing agents and are also stable to hydrolysis. This is important
because it allows reactions to be carried out on substitutions, often under
quite drastic conditions, without destroying the cage structure, rather as
297
the chemistry of derivatives of an aromatic ring such as benzene can be
developed without destroying the ring.
Most chemical studies have been concerned with substituents on
the two carbon atoms. These may be introduced in the first place by
employing substituted acetylene in the carborane syntheses. Such groups
as C-alkyl, -haloalkyl, -aryl, -alkaenyl and -alkenyl may be introduced
into the structure in this way. Further reactions on the subsequents groups
may then be carried out by the usual synthetic methods of organic
chemistry to give, for example, carboxylic acid, ester, alcohol, ketone,
amine or unsaturated groups in the side chain.
The nido-carborane 2.3-C2B4H8 is converted to the closocarboranes C2B3H5, C2B4H6 and C2B5H7 on pyrolysis or ultraviolet
irradiation.
Largely because of preparative difficulties, relatively little is
known about the reactions of the smaller nido-carboranes. They are only
moderately stable to heat and are less resistant to hydrolysis and
oxidation in air than the closo species. Halogen substitutions have been
observed, as has the formation of anions; for example,
+
diglyme
 Na C2B4H7 + H2
C2B4H8 + NaH 
Similarly with LiC4H9, Lithium derivative is former:
 B10C2H10Li2 + 2C4H10
B10C2H12 + 2LiC4H9 
The Sodium derivative with FeCl3 gives Fe-derivative:
 2NaCl + Na2[Fe(B9B2C11)2]
2Na2[B9C2H11] + FeCl3 
298
9.3.3 Structures
Structural studies of carboranes have been done using X-ray
analysis and nmr studies.
The C2B3, C2B4 and C2B5 closo-carboranes, for example, have
trigonal bipyramidal, octahedral and pentagonal bipyramidal skeletal
structrues respectively, and positional isomers have been identified.
The icosahedra structure is similar to that of B 12H122- (Fig. 9.8) and
is electron-deficient, with electron delocalization extending over the
whole framework. It is thus in effect a three-dimensional aromatic
molecule, with marked electron withdrawing character, the most
important result of which is to render the two hydrogen atoms bonded to
carbon acidic. All the C-H and B-H bonds are of the normal two-electron
type and the electron deficiency is associated with the framework, in
which there are multicentre bonds.
The Structure of nido C3B3H7 is shown in Fig. 9.9. In the diagram
hydrogen bridges are shown by curved lines, but terminal B-H and C-H
bonds are ommitted. It can be seen that the introductions of successive
carbon atoms to the framework involves the elimination of one bridge
hydrogen atom and one B-H (i.e. the replacement of BH2 by an
isoelectronic CH unit). Like all the carboranes these compounds are
electron-deficient, with multicentered bonds and delocalization extending
over the entire framework. In much the same way, C2B3H7 has a square
pyramidal structure that is formally derived from that of B 5H9, with two
BH2 replaced by 2CH.
Fig. 9.9
299
9.4
METALLO-BORANES AND METALLO CARBORANES
Borane-clusters, in which metals are present are know as
'Metalloboranes'. Many metalloboranes have been prepared. In some
cases metal atom is attached with the borohydride ion through hydrogen
bridge. The most common and important metalloborane group is one in
which direct metal boron bond is present.
An important example of main group element metallocarborane is
closo [B11H11AlCH3]2- (Fig. 9.10). It is prepared by the action of trim
ethyl aluminium [Al(CH3)3]2 with Na2[B11H13]:

Al2(CH3)6 + 2[B11H13]2- 
2[B11H11AlCH3]2- + 4CH4
Fig. 9.10: Closo [B11H11AlCH3]2The hydrogen attached with carbon in closo- B10C2H12 is slightly
acidic. This can be substituted by butyl lithium or Grignard's reagent to
get lithium or magnesium metallocarboranes:
2C4H9Li + C2H2B10H10  C2Li2B10H10
2RMgBr + C2H2B10H10  [CMgBr]2B10H10 + 2R-H
Similarly, [C2B9H11]2- ion, reacts with FeCl2, BrRe(CO)5 or
BrMn(CO)5 to give Fe, Re or Mn derivatives:
300
2[C2B9H11]2- + FeCl2  [(C2B9H11)2Fe]2- + 2Cl[C2B9H11]2- + BrRe(CO)5  [C2B9H11.Re(CO)3]- + Br- + 2CO
[C2BgH11]2- + BrMn(CO)5  [C2B9H11.Mn(CO)3]- + Br- + 2CO
There is a similar reaction with the hexacarbonyls of Cr, Mo and W
under the influence of ultraviolet light, and the air sensitive products are
of the type (C2B9H11)M(CO)32- (M = Cr, Mo, W). Closely related
complexes of other transition metals (Co, Ni, Pd, Cu and Au) have also
been made, including some with sub-substitutnts on the ion.
In the first place formation of  -bonded complexes based on
carborane structures is not restricted to the C2B9H112- ion; there are a
number formed on the same principle by CB10H113- and some of its
amine-substituted derivatives (e.g. [(CB10H11)2Cr]3- and C2B4H63-) also
give complexes, and it may be noted, some of these are nido-anions. Thus
[1,6 C2B7H9)2Co]- has the structure shown below (Fig. 9.11), the ion
being derived from 1,3-C2B7H13.
(a)
(b)
(c)
Fig. 9.11: (a) Carbonyl metallocene
(b) Carbonyl Cyclopentadieny
(c) Carbolyl Carbonyl Compound
301
On the basis of Wade's rule, the structrues of these metal
derivatives may be known from their molecular formula and skeletal
electrons. For example in B3H7[Fe(CO)3]2, n=5 (3B + 2Fe) and skeletel
electrons are 14. Hence it has nido structure corresponding to square
pyramidal (Fig. 9.12).
Fig. 9.12: Structure of [Fe(CO3)B4H8]
9.4.1 Properties
Just as the carboranes, lithio and Grignard's derivatives of metallo
carbones give substitution reactions of organometallics, which include:
(a) Formation of derivatives such as carboxylic acids, ester, alcohol,
ketone, amines etc.
(b) Synthesis of iodo and nitroso devivatives.
(c) Elmination of Lithium halidePCl3 + C2PhL2B10H10  (C2PhB10H10)2Pl
Ph3PAuCl + C2RLiB10H10  Ph3AuC(Cr)B10H10l
N ( CO) i
2(C6H5)2PCl + C2Li2B10H10  (C6H5)2PC-CP(C6H5)2 

l l
- B10H10 4
OC
CO
Ni
(C6H5)2 PC
B10H10
302
CP (C6H5)2
Similarly, derivatives of mercury and other metals(  -bonded) have also
been obtained,
H Cl
RC2LiB10H10 
 B10H10RC2HgC2RB10H10
9
2
Ph3PAuCl + C2RLiB10H10  Ph3PAuC(CR)B10H10
9.5
METAL CARBONYL AND HALIDE CLUSTERS
As has been described earlier, metal carbonyl clusters are rarely
formed by earlier d-block metals; while that of f-metals are unknown, i.e.
these clusters are formed by group 6 to 10 elements.
An alternative method for counting skeletal electrons in these
compounds is due to D.M.P. Mingos and J. Lauher. This method is also
based on Wade's rule and is known as Wade-Mingos-Lauher rule. In this
method the total number of valence electrons in all the metal atoms
present in the complex are counted and then electrons donated by ligands
are added. Thus in Rh6(CO)166Rh
=
6x9
=
54 e-
16CO
=
16 x 2
=
32 e-
Total =
86 e-
Out of the total 86e-, twelve electrons per rhodium atom are used
for non framework bonding, and remaining 14e- are obtained for skeletal
bonding. These include seven bonding paris, equal to 2n+2 electron.
Hence, Rh6(CO)16 should have closo- structure
Some examples showing inter-relation between cluster-valency
electrons and structures are given in Table 9.2.
303
Table 9.2
No. of Geometry
Metal
Atoms
1.
Monomer
2.
Metal
Bonding
No. of
Skeleton Molecular Cluster
Structure Orbital electron
9
18
Dimer
17
34
Examples
Ni(CO)4
Fe(CO)9,
Mn2(CO)10
3.
Triangle
24
48
Os3(CO)12,
Co3(CO)9CH
4.
Tetrahedron
30
60
Co4(CO)12,
Rh4(CO)12
Butterfly
31
62
Re4(CO)162-,
[Fe4(CO)12C]2-
Square
32
64
Os4(CO)16,
Pt4(O2CMe)8
5.
6.
TBP
36
72
Os5(CO)16
Octahedral
37
74
Fe5(CO)15C
Trigonal
prism
43
86
Ru6(CO)17C
45
90
[Rh6(CO)15C]3-
7.
It is quite clear from table 9.2 in tetranuclear metal cluster three
structures, tetrahedral, butterfly and square planar, are seen, with 60, 62
and 64 cluster electrons respectively (Fig. 9.13).
304
Tetrahedron
Butterfly
Square Planar
Fig. 9.13
Synthesis:
1. Pyrolytic Synthesis:

2CO2(CO)8 
CO4(CO)12 + 4CO 
2. Redox Condensation:
2Ni(CO)4 + [Ni5(CO)12]2- 
 [Ni6(CO)12] + 4CO 
3. Ston's Method: Condensation of meatl carbonyls with unsaturated
metal carbonyls:
(CO)5Mo = C(OMe)Ph + Pt(Cod)2  (CO)5Mo.Pt(Cod)(OMe)Ph
Cp.W(CO)2 (  C.tol) + Co2(CO)8 
 (Cp)(CO)2W.Co2(CO)6C.tol
Reactions:
1. Substitution Fragmentation:
 Fe3(CO)11(Ph3) + Fe3(CO)10(P.Ph3)2 +
Fe3(CO)12 + P.Ph3 
Fe(CO)5 + Fe(CO)4(P.Ph3) + Fe(CO)3(PPh3)2 + CO
2. Prolongation:
[Fe3(CO)11]2- + H+ 
 [Fe3H(CO)11]
3. Cluster Catalytic Ligand Transformation: eg. in Os cluster.
305
9.5
Metal Halide clusters
Although the first information of metal halide clusters was given in
12th Century in the form of calomel, but dimeric nature of mercurous ion
could be established in 20th Century only. But now number of metal
halide clusters are known.
Dinuclear Complexes:
Most important dinuclear species is [Re2X8]2- (Fig. 9.14).
Fig. 9.14: Structure of [Re2Cl8]2Analogous to [Re2X8]2- ion, in which very small M-M distance and
eclipsed configuration of chlorine atoms are present, is [Mo2Cl8]2- and
[W2Cl9]3- (Fig. 9.15).
Fig. 9.15: Structure of [W2Cl9]3-
306
The structures of these dinuclear complexes are either similar to
ethane or an edge-shared bioctahedron or a face shared bioctahedron (Fig.
9.15). or tetragonal prism (Fig. 9.14).
Trinuclear Cluster:
The well known examples of trianuclear cluster are rhenium
trichloride, [ReCl3]3 or Fe3Cl9 and their derivatives. Rhenium Chloride is
a trimer, and has been used for the preparation of other trimers as a
starting material. Its structure is shown in Fig. 9.16.
Fig. 9.16: Structure of [W2Cl9]3-
Tetra nuclear Clusters:
Only a few examples of tetranuclear clusters of halides and oxides
are known. Most important example is the dimeric [Mo 2Cl8]4- cluster
giving a tetra nuclear molecule :
307
Hexanuclear clusters:
Hexanuclear Clusters or Mo, Nb and Ta halides are well known.
Two species are known, one with molecular formula M6X12 or [M6X8]X4
and the other with molecular formula [M6X14]2-.
Molybdenum forms cluster of the type [M6X8]X4. [M6X8]4+ ion has
an octahedral skeleton of metal atoms, each face of which is coordinated
with a chloride ion (Fig. 9.17).
Fig. 9.17: Structure of [M6Cl8]4+
Niobium and tantalum give clusters of M6Cl12 type. In these each
edge of the octahedral structure of metal atoms is coordinated with a
chloride ion (Fig. 9.18).
Fig. 9.18: Structure of [M6X12]
Similarly, Nb, Ta and Zr give clusters of [Nb6Cl12L6]2+ type also. In
which 12 chloride ligands are present (one on each edge) on 12 edges of
the octahedral skeleton of metal atoms and the remaining six ligands are
attached to six metal atoms (one on each metal atom), e.g. Nb and Ta
give [M6X18]2+ type clusters (Fig. 9.19):
308
Fig. 9.19: Structure of [M6X18]2+
Solid [Mo6Cl14]2- species is derived from MoCl2 as its hexamer
(Fig. 9.20).
Fig. 9.20
9.6
COMPOUNDS WITH METAL-METAL MULTIPLE BONDS
As has been shown earlier, the earlier metals in d-block series in
their lower oxidation states have tendency to form metal-metal multiple
bonds. These metal-metal bonds may be present in smaller molecules and
also in macro-chain solids.
309
Chevrel-phases :
Chevrel
phases
generally
involve
tertiary
molybdenum
chalcogenides, MxMoX6, polynuclear clusters, which have characteristic
properties (specially electrical and magnetic). Their structures are also
abnormal. An important example of these phases is a super-conducter
substance, PbMo6S8. Its structure consists of an octahedral cluster of
molybdenum atoms, which is surrounded by cubic cluster of sulphur
atoms. Then this whole structure is enclosed in to a cubic structure of lead
atoms. The internal Mo6S8 cubic structure rotates with respect to lead
lattice. This rotation is due to strong repulsion between sulphur atoms.
Similarly, the superconductivity originates due to overlapping of d-orbital
of molybdenum (Fig. 9.21).
Fig. 9.21  = Mo, o = s, 0 = Pb
Zintle anions and cations:
In 19th century, it was seen that post transition metals in liquid
ammonia solution, in presence of alkali metals give highly coloured
anions. After 1930, polyatomic anions such as Sn94-, Pb74-, Pb94-,Sb73- and
Bi33- were discovered. In 1975 cryptate salts of these anions were also
obtained.
Some cations, Bi95+, Te64+, etc were also prepared
310
These species were designated as Zintle anions and cations. These
are homopoly atomic species, which have no ligands attached with. (Fig.
9.22)
Fig.: 9.22 a. Pb52-, c. Bi95+, d. Te64+
Many compounds having metal metal multiple bonds show ethane
like structure. Important compounds with metal-metal quadruple-bond
include halide complex, [M2X8]2 and oxalate complexes of Cr, Mo and
W. (Fig. 9.23)
Fig. 9.23: structure of [Mo2(CH3COO)4]
311
Table 9.3 gives important informations about the complexes having
metal-metal multiple bonds.
Table 9.3
Complex
Electronic
Configuration
312
Bond Order
Bond
Length
Check Your Progress -2
Notes : (1) Write your answers in the space given below.
(2) Complex your answers with those given at the end of
the unit.
(a) (i) carboranes are......................of carbon and boron having both
these atoms in an...........................skeletal frame work.
(ii) The important example of carboranes is........................which
is isoelectronic with..........................................
(iii) The isomeric compounds of 1, 2 dicarbocloso decabrone 12
are (i)
........................................
and (ii)
........................................
(iv) [C2B9H11]2- ion reacts with FeCl2 and BrRe(CO)5 to give(i)
........................................ and
(ii)
........................................ respectively.
(v) Rh6(CO)16 has total ............................. electrons ...................
electrons per rhodium atom are used for ....................
bonding and remaining ..................... electrons (=
)
pair electrons indicate......................structure.
(vi) In [Nb6Cl12L6]2+ crystals ..................... ligands are present
on ....................... of the ........................... skeleton of metal
atoms and remaining ........................ ligands are attached to
............... atoms.
(vii) Compounds with metal-metal multiple bonds are given by
........................ in .......................... block series, in their
................ oxidation states.
(viii) Example of (i)
Chevrel phase is ........................................
(ii)
Zintle anion is .......................... and ....................
(iii)
Oxalate complex is ........................................
313
9.7
LET US SUM UP
 Closed polyhedrons play important part in the synthesis of cluster
molecules in inorganic chemistry. These cluster molecules include
polyhedral boranes, carboranes, metallo-boroanes and -carboranes
and metal halide crystals.
 Metal clusters include those molecular complexes in which metalmetal bonds from triangular or large closed structures.
 Higher boranes (boron hydrides) may be given general formula
[BnHn]2-, BnHn+4 and BnHn+6. They have different shapes; some
resemble nests, some with butterfly and some with spider's web.
 In higher boranes, in addition to 2c,2e and 3c, 2e bonds, B-B 2c,2e
and B-B-B 3c,2e bonds are also present.
 Wade's rule relates the number of electrons in the higher boranes
with their formulae and shapes. According to this rule the building
blocks are BH units (due to sp hybridisation of boron atoms).
 Out of the two sp-hybrids of B, one is used for 2c,2e B-H bonding
and the other one is directed inside as a radial orbital. Remaining
two unhybridised p-orbital of each B atom are placed perpendicular
to the radial orbitals and are known as tangential orbital. These
radial and tangential orbitals combine to form skeleton of the
deltahedron. To fill the bonding molecular orbitals of the skeleton,
necessary number of electrons are obtained from the radial orbitals
of BH units and extra hydrogen s-orbitals. These electrons are
called skeletal electrons.
 On the basis of structures, molecular formula and skeletal
electrons, higher boranes are classified in to four groups:
Closo, [BnHn]2-, with (n+1) skeletal electron pairs,
Nido, [BnHn+4], with (n+2) skeletal electron pairs,
314
Arachno, [BnHn+6], with (n+3) skeletal electron pairs,
and Hypo, [BnHn+8], with (n+4) skeletal electron pairs,
 Removal of one BH unit and 2 electrons from octahedral closed
[B6H6]2- ion and adding four hydrogen atoms gives square
pyramidal nido B5H9 borane. On repeating same process on nido
B5H9, butterfly shaped arachno B4H10 is obtained. Each of these
three boranes has 14 skeletal electrons, but removal of BH unit
gradually results in more and more open structure.
 The bronsted- acidity of boranes increases with their size:
B4H10 < B5H9 < B10H14
 Carboranes are mixed hydrides of C and B having both these atoms
in an electron defficient skeletal framework. They are classified in
to closo and nido-carboranes accordingly.
 Important member of carboranes is B4C2H6 which is isomeric with
[B6H6]2-.
 Properties of Carboranes resemble with those of the corresponding
boranes. Thus 1,2-di-carbo closo dodecarborane -12 is stable in
both air and heat. Its meta and para isomers are 1,7 C2B10H12 and
1,12 C2B10H12 respectively.
 Borane and carborane clusters in which metals are present are
known as Metalloboranes and Metallo carboranes ClosoB10C2H12 reacts with butyl lithium or grignard's reagent to give
lithium and magnesium metallocarbornes.
Similarly [C2B9H11]2- reacts with FeCl2, BrRe(CO)5 or
BrMn(CO)5 to give Fe, Re or Mn derivatives.
 On the basis of wade's rule the structures of these metal derivatives
may be known from their molecular formulae and skeletal
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electrons e.g. B3H7 [Fe(CO)3]2 with n=5 (3B+2Fe) and skeletal
electron 14 is nido-metalloborane with square pyramidal geometry.

Skeletal electrons in metal carbonyl and halide clusters are
counted using Wade-Mingos-Lauher-rule. In this method the total
number of valency electrons in all the metal atoms present in the
complex are counted and then electrons donated by ligands are
added. Thus in Rh6(CO)16 6 Rh = 6 x 9
= 54e-
16 CO = 16 x 2
= 32e-
Total = 86eOut of these 86e-, 12 per Rh atom are used for the non-frame
work bonding and remaining 14e- are used for skeletal-bonding.
These include (n+1) e--pairs; hence it has closo-structure.
 Important examples of metal halide crystals are [ReCl8]2-, [W2Cl9]3, [ReCl3]3 dimeric [Mo2Cl8]4- and [Mo6Cl8]4+.
 In [Nb6Cl12L6] type clusters, 12 Cl- ligands are present on 12 edges
of the octahedral skeleton of metal atoms and remaining 6 ligands
are attached to six metal atoms.
 Compounds with metal-metal multiple bonds may be either chevrel
phases (e.g. PbMo6S8) Zintle anions or cations (e.g. Sn94-, Pb74-,
Bi95+, Te64+ etc) or metal-metal polybonded complexes.
9.8
CHECK YOUR PROGRESS: THE KEY
1(A) (i)
Metal-metal bonds.
Form a triangular or
Large closed structure.
(ii)
(a)
Nest
(b)
Butterfly
316
(iii)
(B) (i)
(ii)
(c)
Spider's web
(a)
2c, 2e bonds
(b)
3c, 2e bonds
(a)
Number of electrons in the molecule.
(b)
Their formula, and
(c)
Shapes.
Called Deltahedron
Which have BH units.
(iii)
(a)
(b)
(c)
2(i)
(ii)
Name
Closo
Nido
Arachno
Formula
[BnHn]2[BnHn+4]
[BnHn+6]
Skeletal electron-pairs
(n+1)
(n+2)
(n+3)
Mixed hydrides in an electron defficient.
C2B10H12
with B12H12
(iii)
(i) 1, 7 C2B10H12 (meta isomer)
(ii) 1, 7 C2B10H12 (para isomer)
(iv)
(i) [(C2B9H11)2Fe]2- and
(ii) [C2B9H11Re(CO)3]-
(v)
86 electrons.
12 electrons for
non frame work bonding
remaining 14 electrons (=n+1) pairs indicate closo structrue.
(vi)
12 ligands
on 12 edges of the
octahedral skeleton
to six metal atoms.
(vii) Earlier metals in d-block
in their lower oxidation states.
(viii) (i)
PbMo6S8
(ii)
Sn94-, Pb74- and Sb73-
(iii)
[Mo2(CH3OO)4]
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