1

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

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UNIT-1 STEREO CHEMISTRY IN MAIN GROUP COMPOUNDS

Structure:

1.0 Introduction

1.1 Objectives

1.2 Valence Shell Electron Pair Repulsion (VSEPR)

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-]

1.3 Walsh Diagrams

1.3.1 Application to triatomic molecules.

1.3.2 Application to penta atomic molecules.

1.4 d - p Bonds.

1.4.1 Phosphorous Group Elements.

1.4.2 Oxygen Group Elements.

1.5 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.

1.6 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.

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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;

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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

Bonding

Electron Pairs

Shape

BeCl2 2 Linear

BeCl3 3 Triangular Planar

SnCl4 4 Regular Tetrahedral

PCl5 5 Trigonal bipyramidal

SF6 6 Regular Octahedral

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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.

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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 No. of Lone

pairs on

central atom

Bond

Angle

Contraction in

bond angle

w.r.t. CH4

CH4 0 109.5o 0

NH3 1 107.5o 2

o

H2O 2 105.5o 4

o

3. 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 (101.5

o) > PCl3 (100

o)

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 (92.2

o)

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.

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(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)

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(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) (XII)

Fig 1.4 : (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

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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 sp

3 hybridised and has 3bp+1 lp. Thus

H3O+ will also be pyramidal with a bond angle of ~107

o. (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 ~107

o

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)).

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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 120

o. 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 axial-

equatorial (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

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(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.

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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 pictoral-

approach to determine the geometry of a molecule considering and

calculating the energies of molecular orbitals of the molecule.

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The basic approach is to calculate the energies of molecular

orbitals for two limiting structures, say linear or bent to 90o 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 180

o. 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 90o

≤ 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).

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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 PA

X and PA

Z are equivalent non-bonding orbitals labelled

2a, and b1 respectively. The orbitals SA

and PA

Y interact with σ1B and σ2

B,

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 pzA

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.

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(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 180

o; 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.

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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 104.5

o.

(Fig. 1.9).

Fig. 1.9

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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

22 21 ug . In CH2, which has two more electrons than BEH2, three of the

molecular orbital must be occupied. In this case, the lowest energy is

achieved if the molecule is angular and has configuration 12

2

2

1

2

1 1121 baa .

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

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four hydrogen atoms take part. Overlapping of these eight orbitals, eight

molecular orbitals are formed, the four bonding (2σg or 2a1 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 bt1

2

1 2 ). 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 t1u 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

*

1 2t )

molecular orbitals formed; while the d-orbitals [α1(dz2), b1(dx

2 - y

2) 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 .2, 1

2

1

bt 2

13a configuration.

As in the distorted geometry, overlapping of 2t1 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.

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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 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.

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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

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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 200-

300 KJ of NO 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 :

_ + SiH3

H3Si = N

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 Disiloxane

(Sp3-hybridised. 0 + 2 lp) (Sp

3-hybridised. 0 + 1 lp)

Addition compound with BF3 No addition compound withBF3

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Methyl isothicyanate Silyl isothiocynate

(one lp on N; Angular) 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 NSF3 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 dה - pה bonding (compare with

S4H4N4):

Tetra Sulphur tetramide Tetra sulphur tetramide fluoride

(Isoelectronic and (Alternate S = N bond)

Isomorphous with S8

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(ii) In diphenyl phosphonitrilic fluoride, there is evidence of ה bonding

in the ring:

Check Your Progress-3

Notes: 1. Write your answers in the space given below.

2. 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) --------------------------------

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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 109.5

o, 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

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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.

26

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

27

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. AB2

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 sp3dz

2 leading to TBP geometry and Sp

3dx

2-y

2 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

28

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 five-

electron-pair class is possible.

For AB6 molecule, octahedral sp3d

2 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 3

4

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.

29

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 te22te

13te

1) 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.

30

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 Is22s

22p

2 ls

22s

12p

3, 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, SnCl2 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

31

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 = 93.8

o AsH3 = 91.8

o SbH3 = 91.3

o

OH2 = 104.5o SH2= 92

o SeH2 = 91

o TeH2 = 89.5

o

that favours reduction in bond angle in these compounds is the

hybridization discussed above. It costs about 600 KJ mol-1

to hybridize

32

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-s-

rich hybrid, or lower promotion energies and poorer overlap for an s-poor

hybrid. (s-character: sp3 (25%)<sp

2(-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.

33

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 s-

character, 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,

34

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. Write your answers in the space given below.

2. 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

35

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)

36

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

37

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

38

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.

39

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).

40

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 105

o. 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

41

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, sp

3, sp

3d and sp

3d

2

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 bond-

angles 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

42

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 CHECK YOUR PROGRESS: THE KEY

1. (a) (i) Linear

(ii) Pyramidal

(iii) Irregular tetrahedral

(iv) Linear

(b) +NO2 >

.NO2 >

-NO2

2. (a) In H2O molecule, eight electrons are distributed such that

the configuration is 2 2

1 , 1 2

u, 1

2ux, 1

2uy, (i.e. 2a

21, 2b

22,

3a21, 1b

21). Thus the energy of 1 ux (3a1) orbital is very

43

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. (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.

4. (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.

44

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

45

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

46

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 .

47

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

48

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 C2O4

2-.

I-<Br

-<Cl

-<F

-<OH

-<C2O4

2-<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.

49

(3) ESR , NQR and NMR spectra

ESR and NMR spectra of various complexes (IrCl62-

, PtX42-

, PdX42-

etc.) and NQR (Nuclear Quadrapol Resonance) spectra of CoCl42-

clearly 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.

50

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) 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

51

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

2+ B

2 + 2 A B

it is clear that 2 is greater than A

2 + B

2 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

52

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 = A

2 + B

2 - 2 A B

Here ( *)2

is less than A2 + B

2 . 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

53

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

B2) 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

54

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 metal-

complexes can be visualise ; e.g. the formation of bonding and

antibonding molecular orbitals as a result of overlapping of ns orbital of a

55

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 energy-

level 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 LM -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.

56

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 3dz

2 .

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 σ-

orbitals 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

57

approaching towards the central metal cation to form six metal-

ligand 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 non-

bonding 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, 4px, 4py,

4pz, 3dz2 and 3dx

2-y

2, 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:

zzyyxa x-ó

6

1 ...............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:

zxa -2

1 ..................Eg or eg

58

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:

yyy -

2

1 ..................Eg or eg

zzz -2

1 ..................Eg or eg

(c) Since one opposite pair of lobes of 3dz2-y

2 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

yyxxyx

2

122 ..................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-r

2. The proper σ-orbital combination is easily

written down by substituting x2+y

2+z

2 for r

2 in 3z

2-r

2. Thus 3z

2-

r2=3z

2-(x

2+y

2+z

2)=2z

2-x

2-y

2. Consequently the proper combination

for 3dz2 orbitals is:

2 2 yyxxzz yyxxzz 2

Thus,

yyxxzz z-ó22

32

12 ..................T1u or t1u

(a) Overlap of metal 4s-orbital (a1g symmetry orbital) with a group

ligand σ-orbitals (a1g symmetry).

59

zzyyx-xa

σσσσóσ6

12

(b) Overlap of metal 4pz-orbital (t1u symmetry orbital) with a group

ligand σ-orbitals (t1u symmetry).

zzx -2

12

(c) Overlap of metal 3dz2-orbital (eg symmetry orbital) with a group

ligand σ-orbitals (eg symmetry).

yyxxzzz 2232

12

60

(d) Overlap of metal 3dz2-y

2orbital (eg symmetry orbital) with a group

ligand σ-orbitals (eg symmetry).

yyxxyx

σσσσ2

122

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 3dz

2 overlap with six group

ligand σ-orbital viz a, x, y, z, x2

-y2 and z

2 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 z

2 (both with t1u symmetry) overlap to form one σz

2.b

and one σz2*

-MO.

(f) 3dx2-y

2 and x

2-y

2 (both with tu symmetry) overlap to form one σ

x2-y

2b-MO one σ x

2-y

2*-MO.

61

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 a1g 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.

62

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 < T2g < Eg* < A*lg < T1u*

Bonding MO's Non-bonding MO's 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.

63

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. σs

b, σx

2, σy

2, σz

b σx

2-y

2

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 t2g 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*(σx

2-y

2* and σz

2*

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 σx2

-

y2* and σz

2* 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 non-

bonding 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

64

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(x

2-y

2)

2 = (σz

2)

2, (3dx)

2 = (3dyz)

1 =

(3dxz)1 = σ*(x

2-y

2)

1 = σ*(z

2)

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(x

2-y

2)

2 = (σz

2)

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 d-

orbitals, three, the t2g (dxy, dxz and dyz) orbitals are also suitable for the σ-

bonding. 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).

65

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.

66

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 a1g (dz2), a doubly

degenerate eg (dxy and dyz) and the remaining two mono-degenerate

b2g(dxy) and b1g(dx2-y

2) 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)

67

Fig. 2.13: Molecular Orbital Energy Level Diagram in a

Square Planar Complex.

68

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 σ-

bonding 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

69

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 d-

bonding 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:

70

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

71

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.

72

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 p-

orbital 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 p-

orbitals 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

73

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.

74

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 Combination of the ligand

orbitals (LGOs)

t1u px 64522

1yxxy

py 63512

1xyyx

pz 43212

1yxxy

t2g pzx 63152

1yxyx

dyz 46522

1yxyx

dxy 43212

1yxyx

4. 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 ( gt2 ):

75

Bonding -molecular orbitals: t2g = T ggt t22

Antibonding -molecular orbitals: *t2g = T ggt t22

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.

76

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

--, OH

-

etc. 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 eg 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 eg-

molecular 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.

77

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 t2g-

orbitals 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 (1

o<1

o). 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 eg-

MOS. 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

78

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 (LM) -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

79

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 d-

orbitals (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 ML -bond giving large

values of o, hence placed at the top of the series; but the saturated

ligands like F- form LM -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).

80

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?

(i) ...........................................

(ii) ...........................................

(iii) ...........................................

(b) 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.

81

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.

82

Metal complexes may have either ML -bonding or LM -

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 ML -bonding in CN- complexes increases

the value of E, Hence it is at the top of spectrochemical series

(SCS); while LM -bonding in F- complexes very much

reduces E, placing F- at the bottom in SCS.

2. (a)

Metal σ -orbitals Symmetry LGO

(i) Octahedral s and p; dz2 and dx

2-y

2 a1g and t1u ,eg Ʃa, Ʃt1u

and Ʃeg

(ii) Tetrahedral s, dxy, dyz and dxz a1g, t2g Ʃa, Ʃt2g

(iii) Square

planar

dx2

-y2, dz

2, dxy, dyz

and dxz

b1g, g1 , b2g,

eg

Ʃb1, Ʃay

Ʃb2 Ʃeg

83

(b) (i) Eg* - t2g

(ii) t2* - Eg

(iii) B1g* - A1g

3.(a) (i) ML -bonding

(ii) It increase Crystal Field Splitting Energy E very much.

(iii) [Fe(CN)6]4-

(b)

(i) Names (ii) Symmetry (iii) LGO

dxy

t2g

43212

1yxyx

dxz 6315

2

1xxyx

dyz 4652

2

1yxyx

84

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 their stereochemical

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

85

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 :

86

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

87

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

88

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 t2g and eg (or

eg*) orbitals. An s orbital, being spherically symmetrical and non-

degenerate, 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).

89

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 d

n configuration

electron – electron interaction starts.

90

Table 3.1 Splitting of the Terms in a dn ion in Weak Octahedral Field

Configuration

of free ion

Ground

state of

free ion

Energy

level

diagram

Configuration

of free ion

Ground

state of

the ion

Energy

level

diagram

d1

d2

d3

d4

d5

2D

3F

4F

5D

6S

A

B

inverted B

inverted A

no splitting

d6

d7

d8

d9

d10

5D

4F

3F

2D

1S

A

B

inverted B

inverted A

no spliting

* 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 Components in octahedral field

S

P

D

F

G

H

I

A1g

T1g

Eg + T2g

A2g + T2g + T1g

A1g + Eg + T1g + T2g

Eg + Tg + T1g + T2g

A1g + A2g + Eg + T1g + T2g + T2g

91

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 Ground State

Term

Higher Energy Terms

d1, d

2

d2, d

8

d3, d

7

d4, d

6

d5

2D

3F

4F

5D

6S

3P,

1G,

1D,

1S

4P,

2H,

2G,

3F, 2 x

2D,

2P

3H,

3G, 2 x

3F,

3D, 2 x

3P,

1I, 2 x

1G,

1F,

2x

1D,

2x

1S

4G,

4F,

4D,

4P,

2I,

2H, 2 x

2G, 2 x

2F,

3 x 2D,

2P,

2S

92

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

93

of unpaired electrons must not be changed, but this is not quite the same

thing. For example, the transition S2 S

1P

1 is spin-allowed so long as

the spins of the two electrons in the S1P

1 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 (Mn

2+, Fe

3+) 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).

94

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.

95

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 :

dn ion Ground Term Splitting Components

d1 ------------- -------------

d3 ------------- -------------

d5 ------------- -------------

(b) 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.

96

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 dn 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

97

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 t2 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, d

5, d

6 and d

7 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).

98

Dq/B (a)

Dq/B (b)

Fig. 3.6 Tanabe-Sugano diagrams for (a) d2 and (b) d

5 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.

99

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. [CoCl4]2-

is blue

whereas [Co(NH3)6]2+

pink. Multiplicity forbidden transition bands are

sharp, whereas the multiplicity allowed t2geg 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 spin-

multiplicity, but for the terms of different multiplicity, both B and C

parameters become necessary. For the energy difference between the

100

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 d2, d

3, d

7 and d

8

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 d

6 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 d

6 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 d

6

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

101

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:

oB

B

freiontheBinofValue

ComplextheBinofValue

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:-

37.121550

296001

1

1

1

1

1

2

1

1

1

cm

cm

TA

TA

gg

gg

102

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,

3821550 1

11

11

B

cm

B

TA gg

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

).

103

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 LM 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 ML 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

).

104

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 d

5 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

105

lengths. These indicate LM 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

106

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.

107

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 r 1 , (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

108

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.

109

(a) (i) LM charge transfer bands arise due to transfer of charge

from..............................to................................and are exemplified

by.............................complexes.

(ii) ML 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.

110

(ii) LM 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) ML 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

2D d

6

5D

d2

3F d

7

4F

d3

4F d

8

3F

d4

5D d

9

2D

d5

6S d

10

1S

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

111

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:

oB

B

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:

XB

transitionenergyLowest

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

112

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 Splitting Components

d1

2d Eg and T2g

d3

4F A2g, T2g and T1g

d5

6S A1g (No Splitting)

(b)

1st rule Spin multiplicity Spin free oh complexes of Mn

2+(d

5)

2nd

rule Redistribution of

electrons in a

given sub-shell

NiCl42-

, Th complex gives more

intense absorption than of an Oh

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. oB

B

3(a)(i) LM 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.

113

(ii) ML 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.

114

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

115

UNIT-4 MAGNETIC PROPERTIES OF

TRANSITION METAL COMPLEXE

Structure

4.0 Introduction

4.1 Objectives

4.2 Magnetic Moments

4.2.1 Number of Unpaired Electrons

4.2.2 Spin Only Formula

4.3 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

116

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.

117

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

d

x K d

Mx K

M

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

118

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;

kT

Nx corr

M3

2

where corr

MX 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 temperature-

independent 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;

T

X corr

M

2125.0 or TX corr

M83.2

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

119

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:

RT

NxM

3

22

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:

2/1

2/1

2)(84.2

3Tx

N

RTXM

m

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, orbital-

orbital, 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:

2/1)]1([ JJg

120

where J is the total angular momentum quantum number and g is

the Lande splitting factor for the electron, defined as:

)1(2

)1()1()1(1

JJ

LLSSJJg

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;

2/1)]1()1(4[ LLSS

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, d

4 (high spin), d

6 (low spin), d

7 (low spin)

and

d

8 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;

2/12/1 )]1([2)1(4[ SSSS

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;

2/1)]2([ nn

121

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).

122

Table 4.1: Magnetic properties of some complexes of the first-row

transition metals

Central

metal

No. of d

electrons

High spin complexes Low spin complexes

No. of

unpaired

electrons

µ(expt)

BM

µ(calc)b

BM

No. of

unpaired

electrons

µ(expt)

BM

µ(calc)b

BM

Ti3+

1 1 1.73 1.73 - - -

V4+

1 1 1.68-

1.78

1.73 - - -

V3+

2 2 2.75-

2.85

2.83 - - -

V2+

3 3 3.80-

390

3.88 - - -

Cr3+

3 3 3.70-

3.90

3.88 - - -

Mn4+

3 3 3.8-4.0 3.88 - - -

Cr2+

4 4 4.75-

4.90

4.90 2 3.20-

3.30

2.83

Mn3+

4 4 4.90-

5.00

4.90 2 3.18 2.83

Mn2+

5 5 5.65-

6.10

5.92 1 1.80-

2.10

1.73

Fe3+

5 5 5.70-

6.0

5.92 1 2.0-2.5 1.73

Fe2+

6 4 5.10-

5.70

4.90 0 - -

Co3+

6 4 - 4.90 0 - -

Co2+

7 3 4.30-

5.20

3.88 1 1.8 1.73

Ni3+

7 3 - 3.88 1 1.8-2.0 1.73

Ni2+

8 2 2.80-

3.50

2.83 - - -

Cu2+

9 1 1.70-

2.20

1.73 - - -

123

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, d

2, d

6 and d

7 ion complexes and low

spin d4, d

5 ions in octahedral field. In the tetrahedral field, the high spin

d3, d

4, d

8 and d

9 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 Fe

2+,

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

124

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-, CrO4

2- and low spin Co

3+ 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

125

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;

)(

T

Cx

M

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.

(a) (b) (c)

Fig.4.1 Schematic representations of magnetic dipole

arrangements in (a) paramagnetic, (b) ferromagnetic, and (c) anti-

ferromagnetic materials.

Fe(Phen)2(CNS)2 is an example which shows significant variation in

magnetic moment with temperature. (Fig. 4.2)

Paramagnetism

Ferromagnetism

Antiferromagnetism

126

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.

127

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 Tanabe-

Sugano diagram to these d6 complexes show that near the crossover point

between weak and strong field the difference in energy between the spin-

free (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.

5T2g

E

1A1g

Fig.4.3 Variation in energies of 5T2g and

1A1g terms with increasing

for d6 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.

128

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.

129

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 as-

d

MXor

dX K

M

K

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,

T

X corr

M

2125.0 or TX corr

M83.2

130

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;

2/11)]L(L1)S(S[4μ

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;

2/12/1 1)]2[S(S1)S(S[4μ

By recognizing that S will be related to the number of unpaired

electrons (n) by S = n/2, the above expression is simplified to;

2/12)][n(nμ

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, d

2, d

6 and d

7 ion complexes, and low spin d

4, d

5 ions

in octahedral field.

131

The observed magnetic moments for the metals in t2g 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 anti-

ferromagnetic.

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.

132

4.7 CHECK YOUR PROGRESS: THE KEY

1. (i) d

KMXM

(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, d

2, d

6 and d

7 ion complexes, and low spin d

4 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.

133

UNIT-5 METAL π COMPLEXES

Structure

5.0 Introduction

5.1 Objectives

5.2 Metal Carbonyls

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

5.3 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

5.4 Dinitrogen Complexes

5.4.1 Fixation of Nitrogen

5.5 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

134

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 LM, ϭ-bond, has vacant π-orbitals, it can

accept lone pair of electrons from metal-ion (if present) to form ML, π-

bonds. 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 Mπ-bonding stabilizes, the complex, LM

π-bonding destabilize it. This also explains positions of CN- and F

-

ligands 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:

135

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 OCM 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 example-

nickel forms volatile nonpolar carbonyls, where as alkali and alkaline

earth metals from non-volatile electrovalent carbonyls.

136

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;

4

40)(4 CONiCONi C

137

2. When CO is passed over reduced iron at 108

o-220

o 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

to attain noble

gas

configuration

13 12 11 10 9 8

First transition

series

V(CO)6

blue solid

Cr(CO)6

white solid

(sublimes)

Mn2(CO)10

yellow solid

(m.p.154o)

Fe(CO)5

yellow liquid

(b.p.103o) Fe2

(CO)12 black

solid

(sublimes)

CO2(CO)8

orange solid

(m.p.51o)

CO6(CO)16

black solid

Ni(CO)4

colourless

liquid

(b.p.43o)

Second

transition series

Mo(CO)6

white

solid(subli

mes)

Te2(CO)10

white solid

Ru(CO)5

colourless

liquid (m.p.-

22o)

Ru2(CO)9*

Ru3(CO)12

orange solid

(m.p.-154o)

Rh2(CO)8*

Rh4(CO)12

orange solid

Rh6(CO)16

black solid

Third transition

series

W(CO)6

white

solid(subli

mes)

Re2(CO)10

white solid

(m.p.177o)

Os(CO)5

colourless

liquid

(m.p.15o)

Os2(CO)9

Orange solid

Os2(CO)12

Yellow solid

(m.p.224o)

Ir2(CO)8

yellow solid

Ir4(CO)12

yellow solid

Ir6(CO)16 red

solid

138

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:

5

][

24

][

3 )()( CORuICORuRuI COCO

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.

139

yLiAlHCOCrCl 3 o115 Cr(CO)6 + LiCl + AlCl3

2Rul3 + 10CO + 6Ag eatmpressurC250175 2Ru(CO)5 + 6Agl

2Mnl3 + 10CO + 2Mg eatmpressurC21025 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.

2Fe(CO)5 ..VU Fe2(CO)4 + CO

2Fe2(CO)9 heat Fe(CO)5 + [Fe(CO)4]3 + CO

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.

OsO4 + 9CO C100 Os(CO)5 + 4CO2

50 atm pressure

Re2O7 + 17CO C75 Re2(CO)10 + 7CO2

200 atm

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.

2Fe(CO)5 light.V.U Fe2(CO)9 + CO

2Os(CO)5 lightVU .. Os2(CO)9 + CO

140

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 Colour and shape Melting Point, (oC)

V(CO)6 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-155

o

Re2(CO)10 Colourless crystals Sublime at 140o and decompose at 177

oC

Fe(CO)5 Yellow Liquid B.P. 103oC

Fe2(CO)9 Bronze Mica-like

platelets

Decomposes at 100oC

Fe3(CO)12 Dark green crystals Decomposes at 140oC

CO2(CO)8 Orange crystals 51oC

Ni(CO)4 Colourless Liquid B.P. 43oC

141

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

+Fe

2-(CO)4]

-Na2Co3 + H2O

H-atom in [H+Fe

2-(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

142

4. Action of NO: many carbonyls react with nitric oxide (NO) to

form metal carbonyls nitrosyls. For example:

Fe(CO)5 + 2NO C95 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.

Mn2(CO)10 + H2 eatmpressurC 200,200 2[Mn-(CO)5H

+]

0

Co2(CO)8 + H2 eatmpressurC 200,165 2[Co-(CO)4H

+]

6. Action of heat: Different carbonyls yields different products when

heated for example:

Fe(CO)5 C250 Fe + 5CO

3Fe2(CO)9 C70 3Fe(CO)5 + Fe3(CO)12

Fe3(CO)12 C140 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.

143

iii. It may also be prepared by carbonylating CrCl3 with CO in the

presence of a reducing agent like LiAlH4.

CrCl3 + CO + LiAlH4 eatmpressurC 200,175 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.

Cr(CO)6 + 2Na 3.NHLiq Na

2 [Cr2-

(CO)5]2-

+ CO

Cr(CO)6 34 ./ NHLiqNaBH Na

2 [Cr2-

(CO)10]2-

+ 2CO

5. Substitution reactions: Some CO groups present in Cr (CO)6 can

be replaced by pyridine to get a number of products.

Cr(CO)6 Py Cr(CO)4(Py)2 Py Cr2(CO)7(Py)5 33 )()( PyCOCrpy

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.

144

2. Both may also be prepared by the action of CO at 225o 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 50

oC and boils at 175

oC.

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.

2MnCl2+10CO+2Mg (in diethyl ehter) eatmpressur210,C250 Mn2 (CO)10

+ 2MgI2

2. By carbonylating MnCl2 with CO in presence of (C8H5)2CONa

2MnCl2+10CO+4(C8H5)2CONa .140165 atmC Mn2(CO)10+

4(C6H5)2CO+4NaCl

145

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;

Mn2(CO)6 + 2Na 3.NHLiq 2Na [Mn-(CO)5]

-

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.

Mn2(CO)10 + H2 eatmpressurC 200,200 2[Mn-(CO)5H

+]

0

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

146

Preparation:

i. It can be prepared by the action of CO on iron powder at 200oC

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.

FeI2 + 4CO C200 Fe (CO)5I2

iii. It may also be prepared by the action of CO on FeS at 200oC

and 200 atm. pressure in the presence of copper.

2FeS + 10CO + 2Cu eatmpressurC 200,200 2Fe(CO)5 + Cu2S

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.

2Fe(CO)5 C250 Fe + 5 Co

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

147

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+4NO2Fe(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

148

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.

Fe3(CO)12 C140 3Fe + 12CO

iii. Reaction with Na: Carbylate axion is formed when Fe3(CO)12

reacts with Na metal in Liq. NH3.

Fe3(CO)12 + 6Na 3.NHLiq 3Na2+[Fe

2-(CO)4]

2

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.

149

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 3NHlig 2Na[Co(CO)4]

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]+4

CO

(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.

150

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. = 43

oC

ii. It has no solubility in water but dissolved in organic solvents.

iii. It decomposes at 180o-200

oC in to nickel and CO.

Ni(CO)4 3180 C Ni + 4CO

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

151

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.

152

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

153

(MCO, Thus (MCO) types of bonds were assumed to be

present in metal carbonyls. In the formation of MCO 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 Atomic

Number of

the metal

Number of electron

contributed by CO

groups

E.A.N. Succeeding

inert gas

Cr(CO)6 24 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 = 23e-

Mn = 25 e-

Co = 27 e-

6CO = 12 e-

5CO = 10 e-

4CO= 8 e-

V(CO)6= 35e-

Mn(CO)5= 35e-

Co(CO)4= 35e-

154

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.

12

XX

yXG m

Where 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= 1 x 2 = 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 18-

electron rule as shown below for Co2(CO)8.

Co2(CO)8

2Co = 2 x 9e- = 18e

-

8Co = 2 x 8e- = 16e

-

Co-Co bond = 1x 2e- = 2e

-

Co2(CO)8 = 36e

Electrons on one Co atom = 18e-

155

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 π-

bond.

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 atom-

carbon mono-oxide bonds are readily formed in metal carbonyls. It

156

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 (MCO)

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

(MCO), 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 (MCO)

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

(MCO) 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

- + - +

157

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

158

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 C-

atom 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.

159

2. Structure of Fe(CO)5: The various evidences are:

i. 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 1.842A

o

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 d2sp

3-hybrid orbitals six

molecules of CO donate a lone pair of electrons each to six vacant

hybrid orbitals to form six CrCO -bonds as shown in the Figure

5.5. Therefore Cr(CO)6 molecule is diamagnetic in nature and

octahedral in geometry.

160

When chromium atoms form chromium carbonyl [Cr(CO)6]

the metal atom exhibits d2sp

3 hybridisation. Out of 6 d

2sp

3 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

161

Thus, the bond structure of Cr(CO)6 shows 2 kinds of bonds

between Cr and CO.

i. Simple covalent bonds Cr-C 0 (I)

ii. Double bond 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.

162

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 S-

bond) 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 (FeC co-ordinate bond). The presence of Fe-Fe bond

is supported by the diamagnetic character of Fe2(CO)9 molecule. Fe-

Fe bond distance has been found to be equal to 2.46Ao. The terminal

C-O bond distances from the structure given in figure. (5.7)

163

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

164

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

165

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

166

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 one-

electron donation (to each of two metal atoms) by bridging carbonyls,

through the possible existence of bonds of different strengths makes

167

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.

168

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 (:NO:)

+

169

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 II III IV V

:N

:O:

:N

:O:

D

:N

:O:

D

:N

:O:

:N

:O:-

Links with

bonds

(neutral)

Links with

and

bonds

(cationic)

Links with

dative bonds

(cationic)

Links with

dative bonds

(cationic)

Accepts

electron

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,

170

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 N-

O 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

171

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 Fe

1 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 Ru2(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 ML back π-bonding (Fig. 5.13).

172

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

173

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 22

20

ClCH

M(NO)2Cl2 + 6CO

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

+

174

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-

+ 2OH-

These 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-+3Ba

2++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 co-

ordinated 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-

175

+ 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

below-

M+ (:N: ::O:

+) or M

2- - 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

+ M

2- - 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 d

2sp

3 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:

N π O: N – O -N π O

+:

On this basis resonance structures of NO, the metallic

nitrosyls may be represented as:

M--N π O: M- N π O M π N π O: M

--N-O

Nitric oxide is a paramagnetic molecule with an electron in

176

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]

177

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

178

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.

179

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 EtOH

NaBH4 (Ph3P)3CoH3 3N (Ph3P)3CoH(N2)

[Ru(NH3)5H2O]2+

3N [Ru(NH3)5(N2)]2+

(b) Another method of preparation of dinitrogen complex is the

reaction of coordinated azide:

[Ru(NH3)5Cl]2+

N3-

3

3

NH

HMeSO[Ru(NH3)5N2]

2+

(c) Similarly reaction of (Ph3P)2Ir(CO)Cl with RCON3:

(Ph3P)2Ir(CO)Cl + RCON3 Ir(PPh3)2(CO)(Cl)(N2.NCOR)

EtOHCHCl /3 (Ph3P)2IrCl(N2)

180

(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 N

M – N – N M M – N – N – N M M

N N

Terminal Bridging

Structures of the 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.

181

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. Blue-

green algae, symbiotic bacteria legume)

182

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

183

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)

184

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,

185

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.

[(Ph3P)Co(CO)3]-

ICF32 (Ph3P)Co(CO)3I + I- + C2F6

Triphenylphosphine carbonyl halides of rhodium and iridium may

be prepared by interaction of the metal halide (or a complex halide) and

186

triphenylphosphine in a variety of organic solvents, the solvent serving as

the source of the carbonyl group:

(NH4)2IrCl6 PPh3 (Ph3P)2Ir(CO)CI

IrCl3.3H2O PPh3 (Ph3P)2Ir(CO)CI

The product of these reactions- (Vaska's compound) is a highly

reactive complex.

Vaska's compound is a carbonyl halide; and many triphenyl-

phosphine 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,

(Ph3P)2Ir(CO)Cl 2O O2Ir(PPh3)2(CO)Cl

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 O2

2-. Some of the many other reactions of Vaska's

compounds are shown in Fig. 5.24

Fig. 5.24

187

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....................................

188

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 ML π-

bonding, 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 (ML).

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 be-

189

(a) (b) (c) (d)

:N

:O:

:N

:O:

D

:N

:O:

D

:N

:O:

Links with

bonds

(neutral)

Links with

and bonds

(cationic)

Links with

dative bonds

(cationic)

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. 220000-

240000) is a 22 tetramer, which has 2 Mo atoms, nearly 30 Fe

atoms and nearly 30 mobile sulphide ions. Fe-S group functions as

190

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.

191

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.

(c) At 2125-2000 cm-1

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) (c) (d)

:N

:O:

:N

:O:

D

:N

:O:

D

:N

:O:

Links with

bonds

(neutral)

Links with

and bonds

(cationic)

Links with

dative bonds

(cationic)

Links with

dative bonds

(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

192

Unit - 6 REACTION MECHANISM OF TRANSITION METAL

COMPLEXES-I

Structure

6.0 Introduction.

6.1 Objectives.

6.2 Energy Profile of a Reaction.

6.2.1 Reactivity of metal Complex - Inert and Labile Complexes.

6.2.2 Valence Bond and Crystal Field applications.

6.3 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

193

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

194

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

describe the catalysts form octahedral substitution reactions;

6.2 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

195

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) Redox-

reactions. 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

196

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 sp3d

2 hybridisation.

ii. Inner-Orbital Complexes which involve d2sp

3 hybridisation.

The two d-orbitals involved in sp3d

2 and d

2sp

3 hybridisation are

dx2-y

2 and dz

2 eg set orbitals.

1. Outer-Orbital Octahedral Complexes

Outer-orbital octahedral complexes (sp3d

2 hybridisation) are

generally labile for example the octahedral complexes of

Mn2+

(3d5)Fe

2+, Fe

3+(3d

5) Co

2+ (3d

7) Ni

2+ (3d

8) Cu

2+ (3d

9) and Cr

2+

(3d4) exchange ligands rapidly and hence are labile. This is

197

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 d2sp

3 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.

198

Table 6.1 Distribution of dn-electrons in various t2g orbitals for

labile and inert inner-orbital octahedral complexes

(according to VBT)

Type of the

complex

dn

configu

ration

Distribution of dn electron (shown by

arrows) in t2g orbitals. Electrons shown by

crosses in eg orbitals have been donated by

six ligands to enter d2sp

3 hybrids and are

in opposite spins.

Example of central

metal ions

d s p

t2g eg

xy yz zx x2-y

2 z

2 px py pz

inner

orbital

labile

octahedral

complexes

d0 xx xx xx xx xx xx Sc(+3), Y(+3), rare

earth (+3), Te(+4),

Zr(+4), Hf(+4), Ce(+4),

Th(+4), Nb(+5), Ta(+5),

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 xx xx xx xx xx xx [Cr(CN)6]

4-, Mn(CN)6]

1

Re(+3), Os(+3), Ir(+4)

d5 xx xx xx xx xx xx [Mn(CN)6]

4-, Re(+2),

Fe(CN)6]3-

Ru(+3),

Os(+3), Ir(+4)

d6 xx xx xx xx xx xx [Fe(CN)6]

4-, Ru(+2),

Os(+2), Co(+3)

(except Co Fe34-

Rh (+3), Ir (+3)

199

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) (C.N. = 7)

dn ion Strong ligand field (spin-paired or

low-spin complexes)

Weak ligand field (spin-paired or

low-spin complexes)

Oct.

(C.N.= 6)

pent.bipy.

(C.N.= 7)

E Oct.

(C.N.= 6)

pent.bipy.

(C.N.= 7)

E

d0 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq

d1 4 5.28 0 4 5.28 0

d2 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

200

SN1 association mechanism

Octahedral (oct.) Syware Pyramidal (Squ. pyi)

(C.N. = 6) (C.N. = 5)

dn ion Strong ligand field (spin-paired

or low-spin complexes)

Weak ligand field (spin-paired

or low-spin complexes)

Oct.

(C.N.= 6)

pent.bipy.

(C.N.=5)

E Oct.

(C.N.= 6)

pent.bipy.

(C.N.= 5)

E

d0 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.

201

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, d

1 and d

2

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, d

4, d

5,

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 SN

2 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 d3,

d4, d

5 and d

6 ions is:

Order of inertness : d6 > d

6 > d

4 > d

5

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 > d

3> d

4 > d

5 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 SN

2.

202

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 [3dxy

2, 3yz

2, 3dxz

2, 3d(x

2-y

2),

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 d10

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 < 0.99 < 1.13

203

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 14

CN-,

[Ni(CN)4]2+

+ 414

CN- [Ni(

14CH)4]

2- + 4CN

204

while 6-co-ordinated complexes like [Mn(CN)6]4-

and [Co(CN)6]3-

have 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 5th group into the co-

ordination 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 TS

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 ................................

205

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 Products

(Transition State)

Unstable

The difference in energy between the reactants and the activated

complex is called activation energy.

These reaction involves two process (1) SN1 and SN

2

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:

[LnMX] slow [LnM] + X

[LnM] + y fast [LnMY] + X

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

206

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:

[Ln MX] + Y 1k [Ln MX] Y

[Ln MX] y slow [Ln MY] + X

This leads to

dt

d[Ln MY] =

21

21 ]][[

kk

YMXLkk n

= k[Ln MX][Y]

where

k = 21

21

kk

kk

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,

207

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.

208

(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 CNS

-

along 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, d

6 and d

8 and high spin d

3 and d

8 ions. Hence,

these complexes do not favour the SN1 mechanism for the

substitution.

Table 6.4 Changes in LFSE (in Dq) for Changing a 6-

coordinate Complex to a 5-Coordinate (SP) or a 7-Coordinate

(Pentagonal Bipyramid) species.

209

System High spin Low spin

CN = 5 CN = 7 CN = 5 CN = 7

do, d

10 0.00 0.00 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.

MX5Y )( ZSlow MX5YZ

(C.N.=6) 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

210

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 Fast MX5Z

Unstable seven- -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 = ]][[

].[

5

5

ZLMX

ZLMX

The value of the equilibrium constant, K, for the association

complex, may be obtained using Fuoss-Eagan equation,

K = 3

4 a

3 NAe

-v/RT

where, 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.

211

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 d

6 ions require highest activation

energies, followed by d7, d

8 (Ni

2+complexes are labile due to the

expulsion of ligand by the eg orbitals) and high spin d3 and d

8 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 H2O

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,

212

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 (pH10).

Examples are:-

[CoIII

(NH3)5Cl]2+

+ H2O[CoIII

(NH3)5(H2O)]3+

+ Cl-

[CoIII

(en)2ACl]+ + H2O [Co

III(en)2A(H2O)]

2+ + Cl

-

[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+

+ X-

The 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.

Acid hydrolysis

reaction

213

The rate law for acid hydrolysis at low pH thus becomes

-dt

d[Co(NH3)5X] = kA[Co(NH3)5X]

(If X 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 CH3CHNH3 (CH3)2CNH2

< < <

CH2NH3 CH2NH2 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 SN1

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

214

(b) the decrease in the rate of the exchange reaction

[Co(NH3)5(H2O)]3+

H218

O = [Co(NH3)5(H218

O)]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+

+ Cl

slow

[Co(NH3)4(H2O)]

3+ + Cl

-

6-co-ordinate complex 5-co-ordinate Intermediate

[Co(NH3)4(H2O)]3+

+ OH

fast

2 [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.

215

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 co-

ordination 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++ Cl

-

It involves following two mechanisms.

216

1. SN-2

, Displacement Mechanism

The reaction proceeds as:

[Co(NH3)5Cl]2+

)( OHslow [Co(NH3)5(OH)Cl]+

fast [Co(NH3)5(OH)]2+

+ Cl-

(C.N.=6) (C.N.=7) (C.N.=6)

Rate of Reaction = K[Complex][base]

= K[Co(NH3)5Cl][OH-]

The rate law is - dt

d[Co(NH3)5Cl] = KB[Co(NH3)5Cl][OH

-]

2. 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:

[(NH3)5CoCl]2+

+OH- fast [(NH3)4Co(NH2)Cl]

++H2O (6.1)

[(NH3)4Co(NH2)Cl]+

slow [(NH3)4Co(NH2)]2+

+ Cl- (6.2)

[(NH3)4Co(NH2)]2+

+H2O fast [(NH3)5Co(OH)]2+

(6.3)

217

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 be-

dt

d= [Co(NH3)5OH] =

][][

]][)([

222

21

5321

OHKOHK

OHClNHCoKK

= K[Co(NH3)5Cl][OH]

where, K = ][][ 22

2

21

21

OHKOHK

KK

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 5-

coordinate 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.

218

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 co-

coordinating 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 OH

-

and H2O2 occurs as:

OH- + H2O2 H2O + HO2

-

219

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 SN

1CB 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.

[Co(NH3)5(H2O)]3+

slow [Co(NH3)5]3+

fast [Co(NH3)5X]2+

+ H2O

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

[L5M(H2O] 1k [L5M] + H2O

[L5M] + Y 1k [L5MY]

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

220

concentration of Y at which the rate of replacement of water no longer

depends on the concentration of Y. The rate of formation of [L5MY] 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

dt

d[L5M] = k1[L5M(H2O)] - k-1[L5M][H2O] - k2 [L5M][Y]

According to the steady-state approximation, the concentration of

the very reactive [L5M] remains small and constant during the reactions

(i.e. dt

d[L5M] = 0 at the steady state). Thus,

[L5M] = ][][

]M([L[

221

251

YKOHK

OHK

and

dt

d[L5MY] =

][][

]][M([L

221

2521

YKOHK

YOHKK

if k-1 [H2O] > k2[Y]

dt

d[L5MY] =

1

21

K

KK]][M([L 25 YOH

and a second-order reaction will be observed. On the other hand,

if k2[Y] > k-1[H2O]

dt

d[L5MY] = k1 ]M([L 25 OH

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+

, from carbon a to complex,

221

[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.

222

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 -

dt

d = [Co(NH3)5OH] = .................................

= ...................................

(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.

223

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 - TS, 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 d2sp

3

hybridisation for octahedral complication) are inert while the outer

orbital complexes (using sp3d

2 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, SN-1 or

SN-2 mechanisms.

224

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:

[MX5Y] Y [MX5] Z [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:

[LMX5] Z [L-MX5Z] Y [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 [MX5L] 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).

225

The rate of acid hydrolysis reaction is of first order -

[Co (NH3)5X]2+

+ H2O [Co(NH3)5(H2O)]3+

+ X-

As 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+ + Cl

-

SN-2 mechanism gives rate of the reaction = K[Co(NH3)5Cl]2+

[OH-] and the rate law -

dt

d[Co(NH3)5Cl] = KB[Co(NH3)5Cl][OH

-].

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:

[(NH3)5CoCl]2+

+ OH- Fast [(NH3)4Co(NH2Cl)]

+ + H2O

[(NH3)4Co(NH2)Cl]+ Slow [(NH3)4Co(NH2)]

2+ + Cl

-

[(NH3)4Co(NH2)]

2+ + H2O Fast [(NH3)5Co(OH)]

2+

Thus, the rate determining step is the dissociation of the amido

group. The rate law will be-

dt

d[Co(NH3)5OH] =

][][

]Cl][OH)[Co(NH

222

21

-5321

OHKOHK

KK

= 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

226

This reaction is reverse of acid hydrolysis reaction. The same

second order kinetics would be observed for a unimolecular process:

[Co(NH3)5(H2O)]3+

OH

Slow

2 [Co(NH3)5]

3+

X

Fast [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 - TS

That is .......................should be very high i.e. H should be

negative and TS 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 metal-

ligand bond making process.

(ii) SN-1 is Unimolecular.

SN-2 is bimolecular.

(iii) SN-1, rate = K[MX5Y]

SN-2, rate = K[MX5Y][Z]

227

(b)(i) Rate law is -

-dt

d[Co(NH3)5X] =KA[Co(NH3)5X]

(ii) = ][][

]Cl][OH)[Co(NH

22

2

21

-

5321

OHKOHK

KK

= K ]Cl][OH)[Co(NH -

53

(iii) Without breaking M-L bond.

Involves electrophilic attack of proton on oxygen bonded

with metal.

228

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

229

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

230

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.

231

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.

232

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 14

C.

[*Fe2+

(CN)6]4-

+ [Fe3+

(CN)6]3-

[*Fe3+

(CN)6]3-

+ [Fe2+

(CN)6]4-

Ferrocyanide ion Ferricyanide ion Fe-C bonds Fe-C bonds

Low spin and Low spin and longer shorter

inert inert

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 25

oC 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.

233

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- – MnO4

2- to more than 10

6 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]4-

and [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

234

state such as

[Fe (CN)6]4-

..... M+ ..... [Fe (CN)6]3-

A 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

235

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) Reactions involving large size differences proceed slowly.

7.4 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

236

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

237

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 non-

crossing 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.

238

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

239

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

240

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 eg

2) while Co (III)

complexes are low spin complexes (t2g6 eg

o); and

(iii) after the electron – transfer, the configuration in both the

complexes probably becomes t2g2 eg

1 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]4-

241

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 photo-

chemical reaction. Thus, this can open the path of the production of

hydrogen using solar energy. (Creutz and Sution, 1975).

hv

H2O H2 + ½ O2

[Ru* (bpy)3]2+

242

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.

243

7.6 FACTORS AFFECTING ELECTRON TRANSFER

REACTION RATE

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 an-

atomic 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.

244

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 [(NH3)5 Co(H2O)2+

..... (7.3)

245

[NH3)5 Co(H2O)2+

+ 5H2O [Co(H2O]2+

+ 5 NH3 ...... (7.4)

The intermediate formed in reaction (7.1) dissociates to give a 6-

coordinated (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.

246

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 Br

-

catalysed 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.

247

(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.

248

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.

249

* 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.

250

* 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.

251

* 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.

252

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 Thermodynamic Importance of Stability Constants.

8.3 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.

8.4 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

253

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

254

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.

255

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

M + L ML K1 = ]][[

)(

LM

ML

ML ML2 K2 = ]][[

)( 2

LML

ML

ML2 ML3 K3 = ]][[

)(

2

3

LML

ML

Thus MLn-1 + L MLn Kn = ]][[

)(

1 LML

ML

n

n

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.

M + L 1B ML, = ]][[

)(

LM

ML

M +2L 2B ML2, 2 = 2

2

]][[

)(

LM

ML

Thus M + nL nB MLn, n = nLM

MLn

]][[

)( ................(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.

256

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 log10n 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 = 3]][[

)( 3

LM

ML

On multiplying both numerator and denominator by [ML] [ML2]

and on rearranging we get:

3 = ]][[

]][[

]][[

][

2

2

3

3

MLML

MLML

LM

ML

= ]][[

][

]][[

][

]][[

][

2

32

LML

ML

LML

ML

LM

ML

= K1 x K2 x K3

Thus n = ]][[

][

]][[

][2

LML

ML

LM

ML .............

]][[

][

1 LML

ML

n

n

= K1 x K2..........Kn

or n =

nn

n

nK1

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.

257

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 = R

THS

303.2

/ ......................................................(8.3)

258

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 Ca

2+ 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

259

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

260

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:

261

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. (ML) -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

262

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.

263

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 sp3 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 dsp2 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

264

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

So 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 aquo-

complexes of divalent metal ions in the first transition series at dn

configuration.......................and.........................

(v) The Irving Williams order of stability is..........................................

(vi) Chelate effect is primarily due to ............................ factor.

265

8.4 METHODS OF DETERMINATION OF STABILITY

CONSTANTS.

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

266

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 = nm LM

LnMm

][][

][

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

267

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:

L + H+ Ka HL

+, Ka =

]H][L[

]HL[

Basic Ligand Acid

L + M+ KF ML

+, KF =

]M][L[

]ML[

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

+]

268

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

+] -

]H[Ka

]H[CH

[M+] = Cm - [ML

+]

[L] = ]H[Ka

]H[CH

Thus on putting the values of [ML+], [M

+] and [L] from the above

equation in

K1 = ]L][M[

]ML[

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 metal-

ligand 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 ofn A against pH for a proton-

269

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 =

systeminpresentMofnumberTotal

MmetaltoboundLligandofnumberTotal )()(

or n =

n

oi

n

oi

]MLi[.

]MLi[i.

.....................................(8.5)

which can be written using equation (8.1) as,

n =

n

oi

i

1

n

oi

i

[L]β.

β[L].i.

[ = 1]................................(8.6)

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.

n A =

i

oi

iH

i

i

oi

iH

i

[L]β.

[H]β.i.

[ H

o = 1]........................(8.7)

whereas, pL gives the free ligand exponent and may be defined as.

pL = log .][

1

L

270

(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-

n A = o

n

TL

VV

ENVVTLoY

)(

)()(1

1

....................(8.8)

n =

οTCM.n)V(V

TL)E(N)V(V

1

ο

o

οn1iii

....................(8.9)

Where Vo is the initial volume of the solution, E

o, 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

271

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:

pL- log10

0

0 .

)log

1(

TCMnTCL

anti

jn

on

nH

n

. o

iii

V

VV 0

....................(8.10)

Values of proton-ligand formation constants, K H

1 , K H

2 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 H

1 while the pH

value at n A = 1.5 gives the value of K H

2 and so on.

Similarly, the values of stepwise stability constants of metal-

complexes 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

272

(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 = nmxnnm

nmmn

nmPPC

xnmnPnm

])([

])([)1(111

1

111

where, K = Stability Constant

1/K = Dissociation Constant of the complex.

273

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 = 2

1 )]1)(1([

)21)(1(

xPC

xP

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. non-

equimolecular 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 = ))(( xbxa

X

where, K = Stability Constant

X = Concentration of the complex

274

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 = ))(( 11 xbxa

X

=

))(( 22 xbxa

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 = 2)2)(( xbxa

X

Taking the concentration a1, a2 and b1, b2 for the same absorbance

i.e., the same value of x, we have

K = 2

11 )2)(( xbxa

X

=

2

))(( 22 xbxa

X

The value of x is determined from the relation

4x2(a1-a2+b1-b2)-x(4a1b1-4a2b2+b 2

1 - b 2

2 )+ (a1b12-a2b2

2) = 0

275

i.e. AX2 + BX + C = 0

Where, A = 4(a1 -a2 + b1 - b2)

- B = 4(a1b1 - 4a2b2 + b 2

1 - b 2

2 )

or b1(4a1 + b1) - b2(4a2 + b2)

C = (a1b2

1 - a2b2

2 )

By solving the quadratic equation:

K = A

ACBB

2

)4()( 2

K = A

ACBB

2

)4()( 2

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 = M2L

Taking the concentration a1, a2 and b1, b2 for the same absorbance

i.e., for the same value of X, we have

K = )()2( 1

2

1 xbxa

X

=

)2()2( 2

2

2 bxa

X

or 42(a1-a2+b1-b2)-(4a1b1-4a2b2+a 2

1 - a 2

2 + (a 2

1 b- a 2

2 b2) = 0

i.e. AX2 + BX + C = 0

Where, A = 4(a1 -a2 + b1 - b2)

-B = (4(a1b1 - 4a2b2 + a 2

1 - a 2

2 )

or b1(4a1 + b1) - b2(4a2 + b2)

C = (a 1 b 1 - a 2

1 b 2 )

By solving the quadratic equation, the value of X is determined

K = A

ACBB

2

)4()( 2

or X = A

ACBB

2

)4()( 2

Mushran has modified this method so as to suit for 1:3 complexes.

276

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.

c

mx

extpA

A

][

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

277

Stability constant K = ]][[

][

XM

MX

K = ]C).extpA/ACx[(]C).extpA/ACm[(

]C).extpA/A[(

-

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.

278

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,

279

as : n = K1, K2, K3, ..........................Kn

or

nn

n

nKn1

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/r-

ratio) but is inversely proportional with the size of the metal ion.

ML, -bonding increases it, while LM, -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.

ML, -bonding (complexes with the unsaturated ligands)

increases the stability.

280

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 ]L[

I

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.

281

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

nn

1n

nKβn

(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.

(iii) PL- =

]L[

I

(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.

282

Unit - 9 METAL CLUSTERS

Structure

9.0 Introduction.

9.1 Objectives.

9.2 Boranes and Higher Boranes.

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.

9.3 Carboranes.

9.3.1 Synthesis.

9.3.2 Properties.

9.3.3 Structures.

9.4 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.

283

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

284

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

285

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 B-

B-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.

286

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 [B6H6]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-

287

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

Name Formula Skeletal

Electron Pair Examples

Closo [BnHn]2-

n+1 [B5H5]2-

to [B12H12]2-

Nido [BnHn+4] n+2 B2H6 , B5H9, B6H19

Arachno [BnHn+6] n+3 B4H10 , B5H11

Hypo [BnHn+8] n+4 Only derivatives are known.

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. [B5H5]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

288

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:

289

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

290

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.

291

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:

B5H9 + CH3Cl 3AlCl [CH3B5H8] + HCl

292

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) ......................... ............................... .............................

293

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

294

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 dicarbo-

closo-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

295

c. dicarbo-closo-dodecaborone

9.3.2 Preparation:

I(a) The Small Closo Carboranes (C2BnHn+2 where n = 3 to 5)

B5H9 + C2H2 Co490 1,5 - C2B3H5 + 1,6 - C2B4H6 +2,4 - C2B5H7

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.

B10H14 + 2L 2H B10H12L2 22CR R2L2B10H10 + H2 + 2L

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.

296

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 1,2-

dicarbonido - pentaborane(7), C2B3H7, which is prepared as

follows:

B4H10 + C2H2 Co50 C2B3H7 (3 - 4 % yield)

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).

1,2 - C2H10H12 MeO C2B9H12

1H C2B9H13

When C2B9H13 is heated, the closo-undeca-Borone (11) cage

is formed.

297

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 500oC, it is converted

into 1, 7 isomer i.e. meta or neo isomer; while at 700oC 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

298

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 closo-

carboranes 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,

C2B4H8 + NaH diglyme Na+C2B4H7

- + H2

Similarly with LiC4H9, Lithium derivative is former:

B10C2H12 + 2LiC4H9 B10C2H10Li2 + 2C4H10

The Sodium derivative with FeCl3 gives Fe-derivative:

2Na2[B9C2H11] + FeCl3 2NaCl + Na2[Fe(B9B2C11)2]

299

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 B12H122-

(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 B5H9, with two

BH2 replaced by 2CH.

Fig. 9.9

300

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]2-

The 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:

301

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

302

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 halide-

PCl3 + C2PhL2B10H10 (C2PhB10H10)2Pl

Ph3PAuCl + C2RLiB10H10 Ph3AuC(Cr)B10H10l

2(C6H5)2PCl + C2Li2B10H10 (C6H5)2PC-CP(C6H5)2 iCON 4)(

l l

- B10H10 -

OC CO

Ni

(C6H5)2 PC CP (C6H5)2

B10H10

303

Similarly, derivatives of mercury and other metals( -bonded) have also

been obtained,

RC2LiB10H10 29ClH B10H10RC2HgC2RB10H10

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)16-

6Rh = 6 x 9 = 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.

304

Table 9.2

No. of

Metal

Atoms

Geometry Metal

Skeleton

Structure

Bonding

Molecular

Orbital

No. of

Cluster

electron

Examples

1. Monomer 9 18 Ni(CO)4

2. Dimer 17 34 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. TBP

36 72 Os5(CO)16

Octahedral

37 74 Fe5(CO)15C

6. Trigonal

prism

43 86 Ru6(CO)17C

7.

45 90 [Rh6(CO)15C]3-

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).

305

Tetrahedron Butterfly Square Planar

Fig. 9.13

Synthesis:

1. Pyrolytic Synthesis:

2CO2(CO)8 CO4(CO)12 + 4CO

2. Redox Condensation:

Ni(CO)4 + [Ni5(CO)12]2-

[Ni6(CO)12]2-

+ 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)12 + P.Ph3 Fe3(CO)11(Ph3) + Fe3(CO)10(P.Ph3)2 +

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.

306

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]2-

Analogous 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-

307

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 [Mo2Cl8]4-

cluster

giving a tetra nuclear molecule :

308

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):

309

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.

310

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

311

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]

312

Table 9.3 gives important informations about the complexes having

metal-metal multiple bonds.

Table 9.3

Complex Electronic

Configuration

Bond Order Bond

Length

313

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 ........................................

314

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 metal-

metal 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,

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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 Closo-

B10C2H12 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 = 86e-

Out 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

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(c) Spider's web

(iii) (a) 2c, 2e bonds

(b) 3c, 2e bonds

(B) (i) (a) Number of electrons in the molecule.

(b) Their formula, and

(c) Shapes.

(ii) Called Deltahedron

Which have BH units.

(iii) Name Formula Skeletal electron-pairs

(a) Closo [BnHn]2-

(n+1)

(b) Nido [BnHn+4] (n+2)

(c) Arachno [BnHn+6] (n+3)

2(i) Mixed hydrides in an electron defficient.

(ii) 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]