M.Sc. (Previous) Chemistry

Transcription

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 1

2 UNIT-1 STEREO CHEMISTRY IN MAIN GROUP COMPOUNDS Structure: 1.0 Introduction 1.1 Objectives 1.2 Valence Shell Electron Pair Repulsion (VSEPR) Gillespie Laws Applications Comparison of CH 4, NH 3, H 2 O and H 3 O Comparison of PF 5, SF 4 and [ICl - 2 ] 1.3 Walsh Diagrams Application to triatomic molecules Application to penta atomic molecules. 1.4 d - p Bonds Phosphorous Group Elements Oxygen Group Elements. 1.5 Bent's rule and Energetic of hybridisation III and IV groups halides V and VI groups Hydrides halides Isovalent hybridisation Apicophilicity. 1.6 Some simple reactions of covalently bonded molecules Atomic invevsion Berrypseudo rotation Nucleophillic substitution. 1.7 Let Us Sum Up. 1.8 Check Your Progress: The key. 2

3 1.0 INTRODUCTION The concept of molecular shape is of the greatest importance in inorganic chemistry for not only does it affect the physical properties of the molecule, but it provides hints about how some reactions might occur. Pauling and Slater (1931) in their Valence Bond Theory (VBT) proposed the use of hybrid orbital, by the central atom of the molecule, during bond formation. Thus, with the knowledge of the hybridisation used by the central atom of the molecule, one can predict the shape and also the angles between the bonds of a molecule. However, since then more advanced theories have come into existence. In this chapter we explore some of the consequences of molecular shape in terms of VSEPR theory and refine that concept into the powerful concept of molecular symmetry and the language of group theory, using Walsh diagrams and Bent's Theory, towards the end of the unit. You may recall what you have already studied about the directional property of a covalent bond and the concept of hybridisation of orbital to predict molecular geometry. 1.1 OBJECTIVES The main aim of this Unit is to study stereochemistry in main group compounds. After going through this unit you should be able to: discuss absolute shapes of various molecules; explain anomalies in bond angles present in some molecules; describe stability of molecular shapes in terms of energetics, and discuss simple reactions of covalently bonded molecules on the basis of the principles studied in this unit; 3

4 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) 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 Bonding Compound Electron Pairs BeCl 2 2 Linear Shape BeCl 3 3 Triangular Planar SnCl 4 4 Regular Tetrahedral PCl 5 5 Trigonal bipyramidal SF 6 6 Regular Octahedral 4

5 2. When the central atom in a molecule is surrounded by both, bonding electron pairs as well as by lone pairs, then molecule will not have a regular shape. The geometry of the molecule will be disturbed. This alteration or distortion in shape is due to the alteration in bond angles which arises due to the presence of lone pairs on the central atom. How the presence of lone pairs causes an alteration in bond angles can be explained as follows: At a fixed angle the closer the electric-pairs to the central atom, the greater is the repulsion between them. Since the lone-pair electrons are under the influence of only one positive centre (i.e. nucleus), they are expected to have a greater electron density than the bond-pair electrons which are under the influence of two positive centres. Thus lone pair is much closer to the central atom than the bond pair. Hence it is believed that lone pair will exert more repulsion on any adjacent electron pair than a bond pair will do on the same adjacent electron pair. (lp - lp) > (lp - bp)...(i) (lp = lone pair and bp = bond pair) If the adjacent electron pair is a bond pair, then repulsive force between lone pair and bond pair will be greater than repulsive force between two bond pairs. (lp - bp) > (bp - bp)...(ii) On combining relations (i) and (ii) we get (lp - lp) > (lp - bp) > (bp - bp) Thus the repulsion between two lone pairs is maximum in magnitude, that between a bp and lp is intermediate while that between two bond pairs is the minimum. 5

6 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 CH 4, NH 3 and H 2 O molecules. (Table 1.1) Molecules No. of Lone pairs on central atom Table 1.1 Bond Angle Contraction in bond angle w.r.t. CH 4 CH o 0 NH o 2 o H 2 O o 4 o 3. B-A-B bond angle decreases with the increase in electro negativity of atom B in AB 2 molecule where A is the central atom. Example: Pl 3 (102 o ) > P Br 3 (101.5 o ) > PCl 3 (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: H 2 O (105.5 o ) < H 2 S (92.2 o ) Applications of Gillespie Laws Let us take some examples in support of these laws: (a) AX 2 molecule, which has only two bond-pairs, will be linear: X----A-----X Examples in this groups will be BeCl 2, CaCl 2, CO 2 etc. 6

7 (b) If the molecule is AX 3 (I) or AX 2 with a lone pair of electrons on the central atom A, i.e. AX 2 E (II), then the molecule will be triangular (Fig 1.1): (c) Fig 1.1 : (I) = BCl 3, BF 3 etc. (II) = SO 2, SnCl 2 etc. If the molecule is AX 4 (III) or AX 3 E (IV) or AX 2 E 2, then AX 4 will be tetrahedral; AX 3 E will be pyramidal and AX 2 E 2 will be angular. (Fig. 1.2): (d) Fig 1.2 : (III) = CCl 4, CH 4, SiCl 4, GeCl 4 etc. (IV) = NH 3, PCl 3, As 2 O 3 etc. (V) = H 2 O, SeCl 2, etc. If the molecule is AX 5 (VI) or AX 4 E (VII) or AX 3 E 2 (VIII) or AX 2 E 3 (IX) then AX 5 will be triangular bi pyramidal; AX 4 E will irregular tetrahedral; AX 3 E 2 will be T-shaped,; and AX 2 E 3 will be linear. (Fig. 1.3) (VI) (IX) 7

8 (VIII) (IX) Fig 1.3 : (VI) = PCl 5; (VII) SF 4, TeCl 4 etc. (VIII) = ClF 3, BrF 3 etc. (IX) = X e F 2, ICl 2, or I 3 etc. (e) If the molecule is AX 6 (X) or AX 5 E (XI) or AX 4 E 2 (XII) then AX 6 will be octahedral, AX 5 E will be square pyramidal; and AX 4 E 2 will be square planar. (Fig. 1.4) (X) (XI) (XII) Fig 1.4 : (X) = SF 6, WF 6, etc. (XI) = BrF 5, IF 5 etc. (XII) = ICl - 4, X e F 2 etc Comparison of CH 4, NH 3, H 2 O and H 3 O + In table 1.1 bond angles in CH 4, NH 3 and H 2 O molecules are given. In all these molecules, the central atom (C, N and O respectively) is sp 3 hybridised. But they differ in the number of lone pair (s) present on the central atom, being zero in CH 4, one in NH 3 and two in case of H 2 O. Thus the repulsive force between electron pairs gradually increases in these molecules from CH 4 to H 2 O, resulting in the change of geometry and the bond angles. This CH 4 (Four bond pairs) is tetrahedral with the 8

9 characteristic bond angle of o. NH 3 is pyramidal (Three bond pairs and one lone pair) and has a bond angle of 107 o. While H 2 O is angular (Two bond-pairs and two lone-pairs) and has bond angle of 105 o. The increasing lp-lp repulsion decreases the bond angles from o to ~107 o in NH 3 and ~105.5 o in H 2 O. On comparing H 3 O + with these molecules, we notice that, it resembles with NH 3 molecule. As the central atom of H 3 O + on (Oxygen) is also sp 3 hybridised and has 3bp+1 lp. Thus H 3 O + will also be pyramidal with a bond angle of ~107 o. (Fig 1.5) Fig 1.5 :A- CH 4, Tetrahedral, bond angle o B- NH 4, Pyramidal, bond angle ~107 o C- H 2 O, Angular, bond angle ~105 o D- H 3 O +, Pyramidal, bond angle ~107 o Comparison of PF 5, SF 4, ClF 3 and [ICl - 2 ] A comparison of PF 5, SF 4, ClF 3 and [ICl 2 ] - 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 sp 3 d hybridised, and has 0, 1, 2 and 3 lone pairs (lp) respectively. In PF 5, all the five electron pairs (= 10 electrons) are bond pairs and are housed in the five sp 3 d hybrid orbital; resulting in the trigonal bipyramidal geometry of the molecule. (Fig 1.6 (a)). 9

10 SF 4 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 90 o and one bp directly opposite to itself. While in equatorial position (Fig. 1.6 (b)(ii)) it has two bps at 90 o 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 ClF 3., the two lps may be axial-axial (Fig 1.6(c)(i)), or axialequatorial (Fig 1.6(c)(ii)), or equatorial-equatorial (Fig 1.6(c)(iii)) positions. As the axial position will result in maximum repulsion hence the axial position for the lp is ruled out. Thus the molecule will have T. shaped geometry, according to the Fig. 1.6(c)(ii). Similarly, due to reduced lp-lp repulsions and a larger volume that a lp can occupy at equatorial positions, the [ICl 2 ] - 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) PF 5, Trigonal bipyamidal (I) (II) (b) SF 4, Irregular tetrahedral 10

11 (c) ClF 3, T-Shaped (d) [ICl 2 ] -, Linear Fig. 1.6 : Structures of PF 5, SF 4, ClF 3 and [ICl 2 ] - Limitations of VSEPR Theory 1. This theory is not able to predict the shapes of certain transition element complexes. 2. This theory is unable to explain the shapes of certain molecules with an inert pair of electrons. 3. This theory is unable to explain the shapes of molecules having extensive delocalised -electron system. 4. This theory can not explain the shapes of molecules which have highly polar bonds. 11

12 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) BeCl 2...(iv) ClF 3... (ii) PCl 3...(v) XeF 2... (iii) SF 4...(vi) BrF 5... (b) Arrange the following species in the order of increasing bond-angle, considering the repulsive forces due to lone pair: NO 2, NO 2 and Θ NO 2 > > > WALSH DIAGRAMS We know all the systems want to be in a stable state, and the stable state is one in which it has the minimum possible energy. Same is true for the stereochemistry or the geometry of a molecule. The VSEPR theory considered that the most stable configuration of a molecule is one in which repulsive forces between the valence electron-pairs is minimum. In contrast, the molecular orbital theory (MOT) considers that the stable geometry of a molecule can be determined on the basis of the energy of molecular orbitals formed as a result of linear combination of atomic orbitals (LCAO). In 1953 A.D.Walsh proposed a simple pictoralapproach to determine the geometry of a molecule considering and calculating the energies of molecular orbitals of the molecule. 12

13 The basic approach is to calculate the energies of molecular orbitals for two limiting structures, say linear or bent to 90 o for an AB 2 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 AB 2 triatomic molecule (such as BeH 2 and H 2 O) is illustrated in Fig 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 B 2 A or AB 2 molecule is constructed by considering how the composition and energy of each molecular orbital changes as the bond angle changes from 90 o to 180 o. The diagram is in fact just a more elaborate version of the correlation diagram Application to Triatomic Molecules The coordinate system for the AB 2 molecule is shown in Figure 1.8. The AB 2 molecule has C 2v symmetry when it is bent and, when linear D 2 h symmetry. To simplify notations, however, the linear configuration is considered to be simply an extremum of the C 2v symmetry. Therefore the labels given to the orbitals through the range 90 o θ < 180 o are retained even when θ = 180 o. The symbols used to label the orbitals are derived from the orbital symmetry properties in a systematic way, but a detailed explanation is not given here. For present purposes, these designations may be treated simply as labels. (Fig. 1.7). 13

14 Fig. 1.7 The A atom of AB 2 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 P A X and P A Z are equivalent non-bonding orbitals labelled 2a, and b 1 respectively. The orbitals S A and P A Y interact with σ B 1 and σ B 2, σ orbitals on the B atoms, to form one very strongly bonding orbital, 1a, one less strongly bonding orbital, 1b 2, one less strongly bonding 3a 1 and 3b 2. The ordering of these orbitals and in more detail, the approximate values of their energies can be estimated by an MO calculation. Similarly, A for the bent molecule the MO energies may be estimated. Here only p z 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. AB 2 E 2 ) very exact computations are needed for an unambiguous prediction of structure. 14

15 (Angle θ) Figure 1.8 Orbital Correlation Diagram For AB 2 Triatomic Molecules Where A uses only s and p orbitals Form the approximate diagram (Fig. 1.8) it is seen that an AB 2 molecule (one with no lone pairs) is more stable when linear then when bent. The 1b 2 orbital drops steadily in energy form θ = 90 o to 180 o ; while the energy of the 1a 1 orbitals is fairly insensitive to angle. For an AB 2 E molecule the results are ambiguous, because the trend in the energy of the 2a 1 orbital approximately offsets that of the 1b 2 orbiral. For AB 2 E 2 molecules, the result should be the same as for AB 2 E. Since the energy of b 1 orbirtal is independent of the angle. Thus it is not clear in this approach that AB 2 E 2 molecules should necessarily be bent, but all known ones are. 15

16 The H 2 O Molecule: Because of its unique importance, this molecule has been subjected to more detailed study than any other AB 2 E 2 molecule. A correlation diagram calculated specially for H 2 O is shown in the figure. Although it differs in detail for the general AB 2 E 2 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 AB 2 E 2 molecule ought necessarily to be bent. In the diagram calculated expressly for H 2 O 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 106 o, essentially in accord with the experimental value of o. (Fig. 1.9). Fig

17 BeH 2 - Molecule The simplest AB 2 molecule in Period 2 is the transient gas-phase BeH 2 molecule (BeH 2 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 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, BeH 2 is predicted to be linear and to have configuration g 2 u. In CH 2, which has two more electrons than BEH 2, three of the molecular orbital must be occupied. In this case, the lowest energy is achieved if the molecule is angular and has configuration 11a 2a1 b The principal feature that determines whether or not the molecule is angular is whether the 2a 1 orbital is occupied. This is the orbital that has considerable A-2s character in the angular molecule but not in the linear molecule Application to Penta Atomic Molecules For peta-atomic molecules examples of CH 4 and SF 4 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). CH 4 - Molecule Methane, CH 4, has eight valence electrons. During bonding, for orbitals [a 1 g (2s) and t 1u (2p)] of carbon, and one [a 1g (1s)] orbital of each 17

18 four hydrogen atoms take part. Overlapping of these eight orbitals, eight molecular orbitals are formed, the four bonding (2σg or 2a 1 and 2tσ or 2t 1 ) and the four antibonding (2σg* and 2t*σ). The eight valency electrons of CH 4 molecule are distributed in the four bonding molecular 2 orbitals (Electronic Configuration, 2 2 b 1 t 1 ). In the tetrahedral geometry due to overlapping with the orbitals of hydrogen atom the energy of 2a 1 g and 2t 1u orbitals is considerably reduced. In contrast, in the distorted geometry, comparatively less overlapping of t 1u orbitals with the hydrogen orbitals (as compared to that in the tetrahedral geometry) the energy of 2t 1 molecular orbitals increases. Thus the geometry of CH 4 molecule is symmetrical tetrahedral, rather than a distorted tetrahedral. SF 4 - Molecule In the valence shall of sulphur atom, in SF 4 molecule, in addition to 3s (α 1g ) and 3p (t 1u ) orbitals, 3d (t 2g and eg) orbitals are also present. During the bonding, 2p z orbital of each of the four fluoring atoms take * * part. As a result for bonding (2a 1 and 2t 1 ) and four antibonding (2 ) molecular orbitals formed; while the d-orbitals [α 1 (dz 2 ), b 1 (dx 2 - y 2 ) and t 2 (d xy, d xz, d yz )] are present as non-bonding molecular orbitals. Ten valency electrons of SF 4 molecule remain distributed in the four bonding and one 2 b non-bonding molecular orbitals, resulting in 2 2t. 2 1, t 3a configuration. As in the distorted geometry, overlapping of 2t 1 orbitals is comparatively greater (thus reducing their energy) and the filling in 3α 1 orbital considerably reduce the energy of the system as compared to that in the regular tetrahedral structure. Hence SF 4 molecule has a distorted tetrahedral geometry, rather than a regular tetrahedral structure. Thus we can say 'Walsh Diagrams' are complementary to the VSEPR concept. 1 18

19 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 H 2 O molecule on the basis of a Walsh Diagram for an AB 2 molecule (b) Is any AB 2 molecule, in which A denotes an atom of a period 3 element, expected to be linear? If so, which 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 (R 3 Si) 2 NCH 3, (R 3 Si) 3 N and (H 3 Ge) 3 N, the central NSi 2 C, NSi 3 and NGe 3 skeletons are planar. 2. Many tetrahedral species such as SiO -4 4, PO -3 4, SO -2 4 and ClO - 4 have bond lengths shorter than those predicted from conventional tables of single bond radii. In silicates the Si-O-Si units also show what were considered to be Si-O distance that are "too short" for single bonds. 19

20 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 F 3 F 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 d - p Bonding in Phosphorous Group Elements Phosphine Oxide, R 3 P = 0, presents an important example of the participation of d-atomic orbitals of nonmetallic elements in bonding. Presence of bonding is defected with the help of the evidences, such as reduction in the bond length, increase in the bond strength and the stabilisation of charge distribution. On these grounds, compared to 20

21 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 R 3 P = 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 ( KJ, cf KJ of NO bond) and the smaller P-O bond lengths in phosphoryl compounds 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 : _ + SiH 3 H 3 Si = N SiH 3 (ii) The bond angles in disiloxane, H 3 SiOSiH 3, and silyl isothiocyanate, H 3 Si-N=C=S indicates p - d back-bonding (in comparison to ether and methyl isothiocyanate) : Dimethyl ether Disiloxane (Sp 3 -hybridised lp) (Sp 3 -hybridised lp) Addition compound with BF 3 No addition compound withbf 3 21

22 Methyl isothicyanate (one lp on N; Angular) Silyl isothiocynate lp used in bonding, Linear The argument given against the use of d atomic orbitals in bonding by non-metallic elements is that a very high excitation energy is required for the same. Hence it may be concluded that the use of d-atomic orbitals in bonding by non-metallic elements will be possible only in their higher oxidation states and when they are linked with strong electronegative elements, eg. PF 5,SF 6,OPX 3 etc.: (i) The N - S bond length in N SF 3 indicates. The bond order = 2.7 indicating d - p bonding: (ii) Thiazytrifluoride N - S bond length = pm; Bond order = 2.7 (cf N - S = 174 pm, b.0 = 1, N = S = 154 pm; b.0 = 2) In S 4 N 4 F 4 also there is indication of הd - הp bonding (compare with S 4 H 4 N 4 ): Tetra Sulphur tetramide (Isoelectronic and Isomorphous with S 8 Tetra sulphur tetramide fluoride (Alternate S = N bond) 22

23 (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? S 4 N 4 F 4, S 4 H 4 N 4, N SF 3, H 3 SiOSiH 3, CH 3 N = C=S and Et 3 NO. Ans: (i) (ii) (iii) (b) Which are the evidences in favour of d-p bonding in R 3 PO molecule? Ans: (i) (ii) (iii)

24 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 CH 2 F 2, the F-C-F bond angle is less than 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 pxdx 2 apical orbital in TBP structures is often termed "apicophilicity". It is well illustrated in a series of oxysulfuranes of the type- prepared by Martin and Co-workers. These, as well as related phosphoranes provide interesting insight into certain molecular rearrangements. Bent's rule is also consistent with and may provide alternative rationalization for Gillespie's VSEPR model. Thus the Bent's rule prediction that highly electronegative constituents will 'attract' p character and reduce bond angles is compatible with the reduction in regular 24

25 volume of the bonding pair when held tightly by an electronegative substituents. Strong, s-rich covalent bonds require a larger volume in space to bond. Thus double bonded oxygen, despite the high electro negativity of oxygen, seeks s-rich orbitals because of the shortness and better overlap of the double bond. Again, the explanation, whether in purely s-character terms (bent's rule) or in larger angular volume for a double bond (VSEPR), predicts the correct structure. The mechanism operating behind Bent's rule is not completely clear. One factor favouring increased p character in electronegative substituents is the decreased bond angles of p orbitals and the decreased steric requirements of electronegative substituents. There may also be an optimum strategy of bonding for a molecule in which the character is concentrated in those bond in which the electronegativity difference is small and covalent bonding is important. The p character, if any, is then directed towards bonds to electronegative groups. The latter will result in greater ionic bonding in a situation in which covalent bonding would be low anyway because of electronegativity difference. Some light may be thrown on the workings of Bent's rule by observations of apparent exceptions to it. The rate exceptions to broadly useful rules are unfortunate with respect to the universal applications of those rules. They also have the annoying tendency to be confusing to someone who is encountering the rule for the first time. On the other hand, any such exception or apparent exception is a boon to the research since it almost always provides insight into the mechanism operating behind the rule. 25

26 Consider the cyclic bromophosphate ester. The phosphorus atom is in an approximately tetrahedral environment using four σ bonds of approximately sp 3 character. We should expect the more electronegative oxygen atoms to bond to s-poor orbitals on the phosphorus and the two oxygen atoms in the ring do attract hybridizations of about 20%s. The most electropositive constituent on the phosphorus is the bromine atom and Bent's rule would predict an s-rich orbital, but instead it draws another s-poor orbital on the phosphorus atom is that involved in σ bond to the exocyclic oxygen. This orbital has nearly 40% s-character. The oxygen atom ought to be about as electronegative as the other two, so why the difference? The answer probably lies in the overlap aspect. 1. The large bromine atom has diffuse orbitals that overlap poorly with the relatively small phosphorus atom. Thus, even though the bromine is less electronegative than the oxygen, it probably does not form as strong a covalent bond. 2. The presence of a bond shortens the exocyclic double bond and increases the overlap of the σ orbitals. If molecules respond to increase in overlap by rehybridization in order to profit from it, the increased s-character then becomes reasonable. From this point of view, Bent's rule might be rewarded. The p character tends to concentrate in orbitals with weak covalently (from either electro negativity or overlap considerations), and s-character tends to 26

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 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 AB 2 molecules are linear due to the use of linear sp hybrid orbitals. AB 2 molecule should be equilaterally triangular, while AB 2 E molecule should be angular, due to use of trigonal sp 2 hybrids. AB 4, AB 3 E and AB 2 E 2 molecule should be tetrahedral, pyramidal and angular, respectively, because here sp 3 hybrid orbitals are used. These cases are, of course, very familiar and involve no more than an octet of electrons. For the AB 5, AB 4 E, AB 3 E 2 and AB 2 E 3 molecules the hybrid must now include orbitals in their formation. The hybrid orbitals used must be of the sp 3 dz 2 leading to TBP geometry and Sp 3 dx 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 27

28 know experimentally that AB 5 molecules nearly all have TBP structures, the same arrangement is assumed for the AB 4 E cases, and so on. Even this adhoc assumption does not solve all difficulties, since the position preferred by lone pairs must be decided and there is no simple physical model here (as there was in the VSEPR approach) to guide us. A preference by lone pairs for equatorial positions has to be assumed. With these assumptions, a consistent correlation of all structures in this fiveelectron-pair class is possible. For AB 6 molecule, octahedral sp 3 d 2 hybrids are used. For AB 4 E 2 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 PF 5, SF 6 and Te (OH) 6 but not of equal importance with the contribution of the s and p orbitals. 28

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+p 3 = [2(-1806)]+[3(-981)] = 6555 KJ mol -1 For tetrahedral hybridised phosphours. (3 te 2 2te 1 3te 1 ) the energy will be: Ʃ te = 5x (-1187) = 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 Figure 1.10 The energy difference between the hybridised and unhybridised atom represents the increase in energy of the two electrons in the filled 3s orbital and the decrease in energy of the electrons in the half-filled 3p orbitals. The energetics of hybridisation, together with the principle of good overlap, are important in determining the electronic structure of molecules. 29

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 Is 2 2s 2 2p 2 ls 2 2s 1 2p 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 BCl 3, AlCl 3, GaCl 3, InCl 3, TiCl 3 and CCl 4, SiCl 4, GeCl 4, SnCl 2, PbCl 2, can be explained on the basis of hybridisation energy. The heaviest elements of these groups (Thallium and, Tin and Lead) the stable oxidation-states are two unit less than the maximum oxidation states (3 and 4 respectively) i.e. One in Thallium, TlCl and two in Tin and Lead, SnCl 2 and PbCl 2. Although, in these elements (Compared to the light elements of the group) excitation is easy. It is because, in the heavier elements orbitals are more diffused (Fig 1.12), hence in a volume region, values of Tl or Pb are less, compared to B or Si. This results in lesser overlapping of the central atom orbitals with orbitals of chlorine atoms. Hence, in the compounds of heavy elements bonds are weaker, compared to that of the 30

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 PbCl 2. Figure 1.12 Effect of orbital size on overlapping 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) NH 3 = o PH 3 = 93.8 o AsH 3 = 91.8 o SbH 3 = 91.3 o OH 2 = o SH 2 = 92 o SeH 2 = 91 o TeH 2 = 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 31

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-srich hybrid, or lower promotion energies and poorer overlap for an s-poor hybrid. (s-character: sp 3 (25%)<sp 2 (-33%)<sp(50%)); s-rich means>25%; s-poor means<25%) Isovalent Hybridisation In many tetravalent molecules the bond angle is seen slightly distorted, than the ideal 109 o 28' for example in CH 3 Cl, H-C-H bond is 110 o 20'. 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 CH 3 NH 2, the carbon atom in CH 3 OH uses a hybrid orbital having more s-character to link methyl hydrogen. As a result the C-O bond has more p character as compared to the C-N bond. This is an example of Bent's rule. 32

33 1.5.4 Apicophilicity Good example of the effect of the differences in hybrid bond strengths are shown by the bond lengths in MX n Molecules with both equatorial and axial constituents. (Table 1.3) Table 1.3 r eq (pm) r ax (pm) PF PCl SbCl SF ClF BrF An sp 3 d hybrid orbital set may be considered to be a combination of p z d 2 z hybrids and sp x p y hybrids. The former make two linear hybrid orbitals bonding axially and the latter form the trigonal, equatorial bonds. The sp 2 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, PCl x F 5-x, and the alkylphosphorus fluorides, R x PF 5-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 CH 2 F 2 the F-C-F bond angle is less than o, 33

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 p z d 2 z 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 109 o 28' but H-C-H bond angle is... than ' in CH 2 F 2 molecule. (b) Apicophiliaity is It is well illustrated by SOME SIMPLE REACTIONS OF COVALENTLY BONDED MOLECULES One of the major differences between organic and inorganic chemistry is the relative emphasis placed on structure and reactivity. Structural organic chemistry is relatively simple, as it is based on diagonal, trigonal or tetrahedral carbon. Thus organic chemistry has turned to the various mechanisms of reaction as one of the more exciting 34

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 Atomic Inversion The simplest reaction is seen in a molecule of ammonia. This can undergo the simple inversion of the hydrogen atoms about the nitrogen atom. This is analogous to the inversion of an umbrella in a high wind. One might argue that above equation does not represent a reaction because the product is identical to the reactant and no bonds were formed or broken in the process. Leaving aside, the process illustrated above is of chemical interest and worthy of chemical study. Consider the trisubstituted amines and phophines shown in the figure below. (Chiral amines and phophines) 35

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 NH 3. 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 Berry Pseudo Rotation In PF 5 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 PF 5, the two triogonal 36

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 Nucleophillic Substitution The simplest reaction path for neucleophilic displacement may be illustrated by solvolysis of a chlorodialkylphophine oxide. We would expect the reaction to proceed with inversion of configuration of the phosphorus atom. This is generally observed especially when the entering and leaving groups are highly electronegative and is thus favorably disposed at the axial positions, and when the leaving group is one that is easily displaced. In contrast in some cases when the leaving group is a poor one, it appears as though front 37

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 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. O 2 + hv (below 242 mm) O. + O. The oxygen atoms can then attack oxygen molecules to form ozone. O. + O 2 + m m + O 3 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: 3O 2 + hv 2O 3 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: O 3 + hv ( mm) O 2 + O. Thus the process is repeated. 38

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 PF 5, 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 LET US SUM UP After going through this unit, you would have achieved the objectives stated earlier in this unit. Let us recall what we have discussed so for: VSEPR theory considers that the geometry of a molecule depends upon the number of bonding and nonbonding electron pairs in the central atom. These arrange themselves in such a way that there is a minimum repulsion between them so that the molecule has minimum energy (i.e. maximum stability). 39

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 CH 4 (4bp + Olp) is 109 o 28', in NH 3 (3bp + lp) is 107 o and H 2 O (2bp + 2lp) is 105 o. Similarly. BCl 3 is triangular, but SO 2 is angular; PCl 5 is trigonal bipyramidal, but SF 4 is irregular tetrahedral; ClF 3 is T-shaped, while XeF 4 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 H 2 O is angular and BeH 2 is linear; or why CH 4 is tetrahedral and SF 4 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. R 3 PO, H 3 SiOSiH 3, N SF 3, S 4 N 4 F 4 etc. The formation of d -p bond is common for all the second period elements and is not important for the elements of third and higher periods. The p -d bonding is more favorable than the d -p bonding for higher atoms i.e. atoms of third and higher periods. Bent's rule states that more electronegative substituents prefer hybrid orbitals having less s-character and more electropositive substituents prefer hybrid orbitals having more s-character. Thus in 40

41 CH 2 F 2, the F-C-F bond angle is less than o 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 sp 2, sp 3, sp 3 d and sp 3 d 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, SnCl 2 and PbCl 2. Similarly, the hydrides of group VA and VIA elements are found to have bondangles considerably reduced as one moves down in these groups. As the more electronegative atoms use those hybrid orbitals, which have higher p-character and less electronegative atoms use hybrid orbitals with more s-character the bond angle in many tetravalent molecules is seen slightly distorted than 109 o 28', e.g. in CH 3 Cl, HCH bond is 110 o 20'. This is known as Isovalent hybridization. Apicophilicity is the tendency of more electronegative substituents, in trigonal bipyramidal molecules (e.g. PCl 4 F or PCl 3 F 2 ) to seek out the apical orbitals, e.g. in oxysulfuranes. The common reaction of covalently bonded molecules are atomic inversion, Berrypseudorotation; nucleophilic displacement and free radical mechanism. Atomic inversion is the simple inversion of atoms about the central atom of the molecule. This is analogous to 41

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 PF 5 ). 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) + NO 2 >. NO 2 > - NO 2 2. (a) In H 2 O molecule, eight electrons are distributed such that the configuration is 2 2 1, 1 2 u, 1 2 ux, 1 2 uy, (i.e. 2a 2 1, 2b 2 2, 3a 2 1, 1b 2 1). Thus the energy of 1 ux (3a 1 ) orbital is very 42

43 much less, indicating that the molecule is angular. Since in a linear molecule 3a 1 orbital remains non-bonding (i.e. has very high energy). (b) MgCl 2 molecule. 3. (a) (i) S 4 N 4 F 4 (ii) N SF 3 (iii) H 3 SiOSiH 3 (b) (i) (ii) (iii) The dipole mement has low value (14.6 x 10-3 Cm). The bond dissociation energy is high ( KJ). P-O bond length is short. 4. (a) (i) Monovalency of Thallium (TlCl) and divalency of Lead (PbCl 2 ). (ii) (iii) Hybrids of VA and VI A group elements on moving down the group. FCF bond angle is shorter than 109 o 28', but HCH bond angle is larger than 109 o 28'. (b) Apicophilicity is the tendency of more electro-negative substitutes to seek out apical orbital in a triangular pipyramidal (TBP) structure. It is well illustrated by oxysulfuranes. 43

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) Octahedral complexes Tetrahedral complexes Square Planar complexes similarity in CFT and MOT 2.4 Ligand Field Theory (LFT) 2.5 Let Us Sum Up 2.6 Check Your Progress : The Key 44

45 2.0 INTRODUCTION The accidental discovery of hexaammine cobalt (III) chloride, CoCl 3 6NH 3 by Tassaert in 1799 raised the question how and why CoCl 3 and NH 3 each of which is stable compound, could combine to give yet another very stable compound, CoCl 3-6NH 3. 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 (t 2 g) of lower energy and 45

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. 46

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 47

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 H 2 O and - OH; F - and CN and CO and C 2 O I - <Br - <Cl - <F - <OH - <C 2 O 4 2- <H 2 O<Pyridine<NH 3 <ethylene diammine< dipy <O - Phenan.<CN - <CO. (2) Nephelauxetic effect The interelectronic repulsions as reflected by the actual spectra of the complexes are less than those present in the free ions. This is attributed to the effective increase in the size of the orbitals housing the electrons by the combination of the orbitals of the metal ion and those on the ligands giving larger room to the elctrons to move around increasing, at the same time, the stability of the complexes. This is termed as the nephelauxeticeffect (cloud expansion effect) and is produced by the large sized ligands having d or other suitable orbitals, like * orbitals. 48

49 (3) ESR, NQR and NMR spectra ESR and NMR spectra of various complexes (IrCl 2-6, PtX , PdX 4 2- etc.) and NQR (Nuclear Quadrapol Resonance) spectra of CoCl 4 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 M 2+ O 2- M 2+. Because of the overlap of the opposite spin electrons of the O 2- ion with the two adjacent atoms, the two metal atoms should also have opposite spins. 49

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) (b) Proofs of covalency in complexes are presented by the evidences such as (i) (ii) (iii) The anomaly in the positions of F and CN ions in the spectrochemical series can be explained in terms of. 2.3 MOLECULAR ORBITAL THEORY (MOT) The molecular orbital theory was developed by Hund and Mulliken. According to Molecular orbital theory, which is modern and more rational, a molecule is considered to be quite different from the constituent atoms. All the electrons belonging to the atoms constituting a molecule are considered to be moving along the entire molecule under the influence of all the nuclei. Thus a molecule is supposed to have orbitals of varying energy levels in the same way as an isolated atom has. These orbitals are called molecular orbitals. The molecular orbitals are filled in the same way as the atomic orbitals, following the Aufbau and Pauli's Exclusion Principles. Thus, a molecular orbital, like an atomic orbital, can contain a maximum number of two electrons and the two electrons have opposite spins. The difference between an atomic orbital and a molecular orbital is that while an electron in a atomic orbital is influenced by one positive 50

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 H A and H B. 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 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 51

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 A 2 + 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 H 2 molecule(fig. 2.2). Fig. 2.2: Formation of s bonding molecular orbital The high electron concentration in between the two positive nuclei shields them from mutual repulsion and holds them together at the observed distance from each other. Such a molecular orbital is said to be a bonding molecular orbital. It is designated as (1s) orbital. The sign 52

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. A 2 and B 2 ) are shown by dotted lines. When the probability densities of both the atoms are added, the probability density function for the bonding MO i.e. 2 MO is obtained. It is shown by solid lines. It is clear from the following figure that there is greater probability of finding the electron between the two nuclei than for separate atoms (Fig. 2.4). Probability Density Plot Fig. 2.4 : Sum combination of Wave-function 53

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 H 2 Similar to the formation of bicentric molecular orbital in H 2 molecule, the formation of polycentric molecular orbitals in metalcomplexes can be visualise ; e.g. the formation of bonding and antibonding molecular orbitals as a result of overlapping of ns orbital of a 54

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 energylevel diagram. 5. In the final step the total bonding electrons ( i.e. valency electrons in the meta ion + the n number of electron - pairs obtained from n ligands as a result of LM -bond formation) are distributed in the molecular orbitals according to aufbau and Hund rule Octahedral Complexes In an octahedral complex, the metal ion is placed at the centre of the octahedron and is surrounded by six ligands which reside at the six corners of the octahedra (i.e. along +x, -x, +y, -y, +z and -z axes (Fig. 55

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, 4p x, 4p y, 4p z, 3d xy, 2 3d zx, 3d 2 2 x -v and 3d z. All the nine atomic orbitals have been grouped into four symmetry classes which are given below: 2 2 4s A 1g, or a 1g ; 4p x, 4p y, 4p z T 1u or t 1u, 3d x -y 3d z 2 E g, or e g ; and 3d xy, 3d yz, 3d zx T 2g or t 2g 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 σ x, σ -x. σ y, σ -y, σ z 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, 4p x, 4p y, 4p z, 3dz -2 2 and 3d 2 x -y, since these orbitals have their lobes lying along the axes along which the six σ-orbitals of the six ligands are 56

57 approaching towards the central metal cation to form six metalligand σ-bonds. The remaining three AO's namely 3d xy, 3d yz and 3d zx do not participate in σ-bonding process, since these have their lobes oriented in space between the axes. Thus these orbitals remain nonbonding and hence are called non-bonding orbitals. 2. Construction of LGO for σ-bonding The σ-orbitals of ligands combine together linearly to form such group of ligand σ-orbitals that should be capable of overlapping with the central metal ion six AO's viz 4s, 4p x, 4p y, 4p z, 3d 2 2 z and 3d 2 x -y, 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: σ x + σ -x + σ y + σ -y + σ z +d -z. This linear combination is represented by a which in its normalised form, is given by: a 1 x 6 ó - x y y z z...a 1g or a 1g (A 1g or a 1g represents group symmetry class name of a ) (b) Since one lobe of 4p x orbitals has + sign and the other has - sign, the linear combination, of ligand σ-orbitals that can overlap with 4p z orbital is σ x -σ -x. It is represented by x which in its normalised form is given as: a 1 x - 2 z...e g or e g 57

58 58 Similarly σ y -σ -y and σ z -σ -z are the linear combination of ligand σ-orbitals that overlap with 4p y and 4p z atomic orbitals respectively. Thus: y y y E g or e g z z z E g or e g (c) Since one opposite pair of lobes of 3d z 2 -y 2 orbital has a + sign and the other has - sign, the linear combination of ligand σ-orbitals for this orbitals is: σ x + σ -x - σ y σ -y. Thus y y x x y x E 1u or e 1u (d) To find the ligand σ -orbitals combination for 3d z 2 orbital poses some difficulty. The analytical function for3d z 2 orbitals is proportional to 3z 2 -r 2. The proper σ-orbital combination is easily written down by substituting x 2 +y 2 +z 2 for r 2 in 3z 2 -r 2. Thus 3z 2 - r 2 =3z 2 -(x 2 +y 2 +z 2 )=2z 2 -x 2 -y 2. Consequently the proper combination for 3d z 2 orbitals is: 2 2 y y x x z z y y x x z z 2 Thus, y y x x z z z 2ó T 1u or t 1u (a) Overlap of metal 4s-orbital (a 1g symmetry orbital) with a group ligand σ-orbitals (a 1g symmetry).

59 59 z z y y -x x a σ σ σ σ ó σ (b) Overlap of metal 4p z -orbital (t 1u symmetry orbital) with a group ligand σ-orbitals (t 1u symmetry). z z x (c) Overlap of metal 3d z 2 -orbital (e g symmetry orbital) with a group ligand σ-orbitals (e g symmetry). y y x x z z z

60 (d) Overlap of metal 3d 2 z -y 2 orbital (e g symmetry orbital) with a group ligand σ-orbitals (e g symmetry). 1 σ σ 2 σ σ 2 2 x y x x y y 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 2 cation viz 4s, 4p x, 4p y, 3d 2 x -y and 3d 2 z overlap with six group 2 ligand σ-orbital viz a, x, y, z, 2 x -y and 2 z respectively to form six sigma bonding (abbreviated as σ b ) and six sigma antibonding (abbreviated as σ*) molecular orbitals. Thus- (a) 4s and a which have the same symmetry (a 1g symmetry) overlap b * to form one σ s MO and one σ s MO. b (b) 4p x and x (both with e g symmetry) overlap to form one σ x -MO * and one σ x -MO. b (c) 4p y and y (both with e g symmetry) overlap to form one σ y -MO * and one σ y -MO. b (d) 4p z and z (both with e g symmetry) overlap to generate one σ z - * MO and one σ z -MO. (e) 3d 2 z and 2 z (both with t 1u symmetry) overlap to form one σz 2.b 2* and one σ z -MO. (f) 3d 2 x - 2 y and 2 x - 2 y (both with t u symmetry) overlap to form one σ x 2-2b- y MO one σ x 2-2*- y MO. 60

61 Thus we see that the combinations of six central metal atomic orbitals with six ligand σ-orbitals give six σ b and six σ* molecular orbitals in an octahedral complex. It is thus obvious that on adding twelve molecular orbitals (6σ b - and 6σ * -MO's) to the three non-bonding AO's viz. 3d xy, 3d yz, 3d zx 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): σ ligand < ligand <3d <4s <4p. (ii) The mixing of the metal and the ligand group orbitals is proportional to the overlap of the metal and ligand orbitals, and is inversely proportional to their coulombic energy difference. (iii) Bonding σ MOs are more stable than the bonding orbital and therefore, the antibonding σ*mos are less stable than the antibonding * orbitals. (iv) The bonding orbital lie closer to the LGOs, the antibonding orbitals are closer to the metal orbitals due to the electronegativity differences between the metal ion and the ligands atoms. As the 4s and 4p orbitals can have a better overlap with LGOs than the 3d orbitals can have with the LGOs, the a 1g and t 1u, MOs are at the lowest energy (corresponding a 1g * and t 1u * go to the highest energy levels). The e g and e g * orbitals derived from the e g orbitals of CFT, interact less with the LGOs due to a poorer overlap, while the t 2g orbitals remain nonbonding in a σ only picture and are not displaced. The energy level diagram is given in Fig for the octahedral complexes. 61

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. T 1u, T 2g and T 1*u ). two are doubly degenerated (i.e. Eg and Eg*) and the remaining two are monodegenerated (i.e. A lg and A* lg ). These energy levels are arranged as follows: A 1g < T 1u < E g < T 2g < E g * < A* lg < T 1u * 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. 62

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 σ b -MO's (i.e. σ b s, σ 2 x, σ 2 y, σ b z σ 2 2 x - y and σ 2 z 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 σ b -MO's are mainly localised on ligand orbitals, since σ-orbitals of the ligands are more stable than the metal orbitals. Conversely, electrons occupying any of the six σ * -MO's are to be considered mainly metal ion electrons. Electrons in t 2g set of orbitals will be purery metal ion electrons when there are no ligand orbital. Further, the crystal field splitting energy ( o or 10 Dq) in an octahedral complex, according to MOT, is the difference in energy between the t 2g (d, 2 xy d yz, and d zx orbitals) and e g *(σ 2 x -y * and σ 2 z * MO's) energy levels. In case of weaker ligands such as F-ion, the energy difference, o between the t 2g set e g * set is smaller than P (i.e. 2 o <p) and hence the lowest-energy antibonding MO's namely σ x - y 2 * and σ 2 z * have approximately the same energy as the nonbonding AO's: 3d xy, 3d yz and 3d zx (t 2g set). Consequently spin free (or high spin) complexes are formed. However, the strong(er) ligands such as NH 3 molecules split the σ-bonding MO's more widely and the energy difference, o between the t 2g -set of nonbonding AO's (3d xy, 3d yz and 3d zx ) and e g *-set of MO's (σ 2 x -y * and 2 σ 2 g *MO's) is greater than the electron pairing energy. P (i.e. o > P). Hence spin paired (or low spin) complexes are formed. For example Co(III) complexes in an decahedral stereochemistry, in 63

64 all have 18 valence electrons, 12 electrons from six ligands and 6 from d-orbitals in Co 3+ ion. The distribution of 18 electrons in [CoF 6 ] 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: (σ s ) 2, (σ x ) 2 = (σ y ) 2 = (σ z ) 2 2, σ( 2 x -y ) 2 = (σ 2 z ) 2, (3d x ) 2 = (3d yz ) 1 = (3dxz) 1 = σ*( 2 x - 2 y ) 1 = σ*( 2 z ) 1 Hence it will be paramagnetic due to four unpaired electrons. While, the distribution of electrons, in [Co(NH 3 ) 6 ] 3+ which contains strong(er) ligans viz. NH 3 molecules (i.e. it is a low-spin complex) takes place in various orbitals as follows: (σ s ) 2, (σ x ) 2 = (σ y ) 2 = (σ z ) 2 2, σ( 2 x -y ) 2 = (σ 2 z ) 2, (3d xy ) 2 = (3d yz ) 2 = (3dxz) 2 Thus the complex is diamagnetic, due to all paired electrons 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 (a 1g ) and p (t 1u ). In addition to this, out of the five d- orbitals, three, the t 2g (d xy, d xz and d yz ) orbitals are also suitable for the σ- bonding. Thus the combined symmetry of p(t 1u ) and d(t 2g ) orbitals is generally expressed by T 2 (C f 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 T 2 symmetry, t 2, and one LGO of a 1 symmetry. The LGO of T 2 symmetry will overlap metal ion orbitals of both the groups (p and d, i.e. t 1u and t 2g ). 64

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) e g orbitals are the non-bonding orbitals; and the energy difference between the e g and the next higher t 2 * 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 [CoCl 4 ] 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. t 2, a 1g and e g 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 [CoCl 4 ] 2- complex is paramagnetic due to three unpaired electrons. 65

66 2.3.3 Square Planar Complexes The square planar complexes, ML 4 have D 4 h symmetry. In these complexes five d-orbitals of the metal ion, loosing their degeneracy, split into the four groups of symmetry a mono degenerate a 1g (d 2 z ), a doubly degenerate e g (d xy and d yz ) and the remaining two mono-degenerate b 2g (d xy ) and b 1g (d 2 x - 2 y ) orbitals. Similarly, metal ion p(t 1u ) orbitals also loose their degeneracy and split into two groups, a monodegenerate a 2u (pz) and a doubly degenerate (p x and p y ) symmetry orbitals. The four ligands present on the two x and y axes give ligand group of orbitals (LGO) of a 2g, b 1g and e u 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 a 2u, e g and b 2g symmetry orbitals remain non-bonding in the complex. (Fig. 2.13) 66

67 Fig. 2.13: Molecular Orbital Energy Level Diagram in a Square Planar Complex. 67

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 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 t zg and e g groups. The energy of e g group being higher than that of t 2g group (Fig. 2.1). Similarly, according to MOT, after σ-overlap of metal orbital with ligand group orbitals (LGO) in the octahedral stereochemistry 68

69 the t 2g symmetry orbitals form the non-bonding energy level in the complex, while the anti-bonding e g, E* g, form the higher energy level above the T 2g level. (Fig. 2.10) Further, similar to CFT, in MOT also, the energy difference between t 2g 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) 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, PX 3 (x=halogen), PR 3, AsR 3, SR 2, C 2 H 4 etc. As a matter fact -bonding plays an important part in their stability. - bonding may be of four types: 69

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, NO - 2, 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. PF 3, PR 3, AsR 3, SR 2 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 O 2-, F -, NH 2 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 O 2-, F -, NH - 2 etc. (Fig. 2.17). Fig. 2.14: M(d ) L(p ) bonding 70

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 Formation of - Molecular Orbitals: -Bonding in Octahedral Complexes: Construction of -molecular orbitals follows the first three steps, used in the formation of σ-molecular orbitals; e.g. (i) Selection of metal - orbitals (ii) Construction LGO's of matching symmetry and (iii) formation of -bonding molecular orbitals as a result of the overlap of -metal orbitals with the LGO's. 71

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 t 2g or t 1u symmetry. As the orbitals of t 1u symmetry (i.e. p orbitals) if used for -bonding will weaken - bonding, hence the t 2g 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. PH 3, PR 3, AsR 3 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. t 2g ), 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 L 1 1x, 1y on -x axis In ligand L 2 2x, 2y on y axis In ligand L 3 3z, 3x on -y axis In ligand L 4 4z, 4x on z axis In ligand L 5 5x, 5y on -z axis In ligand L 6 6x, 6y 72

73 In Fig fot Mp -Lp bonding the overlap of p z 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 p z orbital of the metal ion. Similarly, for Md - Lp bonding, overlaps of metal d xz 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 t 1u and t 2g symmetries. 73

74 74 Fig. 2.20: LGOs for d xz metal orbitals. Table 2.1: LGO's matching symmetry with metal t 1u and t 2g group orbitals for bonding is an octahedral complex: Symbol Metal Orbital Combination of the ligand orbitals (LGOs) t 1u p x y x x y p y x y y x p z y x x y t 2g p zx y x y x d yz y x y x d xy y x y x 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 (t 2g orbitals) with the LOGs of same symmetry ( g t 2 ):

75 Bonding -molecular orbitals: t 2g = T Antibonding -molecular orbitals: *t 2g = T t 2 g t 2 g t 2 g t 2 g It is important to note that both the bonding and anti-bonding molecular orbitals will be triply degenerated. 5. Accommodation of molecular Orbitals in the molecular orbital energy level diagram The energy of -molecular orbitals depends on two factors: (i) as compared to the energy of metal ion -orbitals (t 2g ), whether the energy of LOGs t 2g symmetry is lower or higher; and (ii) whether the ligand -orbitals are saturated or vacant? This information is important in the light of the fact that during combination of the orbitals the energy of orbitals having lower energy is further lowered, while that of higher energy is raised further. Hence; (a) when the ligand orbitals are vacant and are of higher energy than that of the filled metal orbitals, M L, -bonding takes place; but (b) when the ligand orbitals are saturated and are of lower energy than that of the metal orbitals, L M, -bonding takes place. Metal ions placed towards the end of a transition series in their lower oxidation state give the first type of -bonding (i.e. M L, -bonding) with the unsaturated ligands, such as phophine, arsine, CN -, CO, NO etc. This type of -bond increases stability of the complex. 75

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 t 2g -orbitals, t 2g -bonding and t 2g * anti-bonding molecular orbitals are accommodated. When there is M L, -bonding, the bonding t 2g molecular orbitals (triply degenerated) find their place above the e g group of orbitals, while the anti-bonding *t 2g orbitals get their position above the anti-bonding e g *, but below the anti-bonding a 1g * orbitals (Fig. 2.21) Similarly, in case of L M, -bonding, the triply degenerate bonding t 2g orbitals are placed above the bonding e g - molecular orbitals, while the anti-bonding - t 2g * orbitals find their place below the anti-bonding e g *-molecular orbitals (Fig. 2.22): Fig. 2.21: M L, -bonding. 76

77 Fig. 2.22: L M, -bonding Effect of -bonding on ligand field splitting energy o As due to -bonding the energy of non-bonding t 2g 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 t 2g - orbitals are stabilised relative to the e g *-MOs i.e. the metal t 2g electrons will go into the t b 2g -MOs which are of lower energy than t 2g *-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 metalto-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 t 2g metal orbitals are de-stabilised relative to e g - MOS. The electrons of the ligand -orbitals enter the lower t * 2g -MOs and those of the t 2g metal orbitals will go to higher t 2g * MOs. Since in this 77

78 case, the -interaction destabilises the t 2g -metal orbitals relative to e * g - MOs, the value of o is diminished to o as shown in Fig ( 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 Ligand field theory: An adjusted crystal field theory (ACFT) As has been pointed out earlier, ligand field theory was specially developed to explain the nature of metal ligand bond in metal complexes. We know, the crystal field theory takes no account of possible covalent bonding in complexes and regards the bonding as purely electrostatic. But the physical measurements such as electron spin resonance. NMR and nuclear quadruple resonance suggest that there is some measure of covalent bonding also in complexes. It is because of this reason that a kind of modified form of CFT has been suggested in which some parameters are empirically adjusted to allow for covalence in complexes without explicitly introducing covalence into CFT. This modified form of CFT is often called 'Ligand Field Theory' (LFT). Ligand field theory, which is a combination of crystal field theory and molecular orbital theory, is therefore in principle the most satisfactory of the theories of bonding discussed in this unit. It may be considered as the combination of -bonding in CFT or we may say LFT calculates the effect of -bonding on the crystal field splitting. Thus, LFT explains the mechanism of metal-ligand bonding in two steps. The 78

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 H 2 O 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). 79

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 LET US SUM UP Although Werner (1893) proposed two types of valencies (Primary and Secondary) for metal ions in, so called the compounds of higher order (i.e. metal complexes) and Sidgwick (1927) introducing the concept of coordinate bond and EAN explained primary valency as electrovalency and the secondary valency as coordinate covalency. The exact nature of M-L bond in the complex compounds could not be understood, until the quantum mechanical theory was used. 80

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. 81

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 t 2g 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 H 2 O 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 2. (a) the value of E, Hence it is at the top of spectrochemical series (SCS); while LM -bonding in F - reduces E, placing F - at the bottom in SCS. (i) Octahedral s and p; d z 2 and d x 2 -y 2 complexes very much Metal σ -orbitals Symmetry LGO a 1g and t 1u,e g Ʃa, Ʃt 1u and Ʃe g (ii) Tetrahedral s, d xy, d yz and d xz a 1g, t 2g Ʃa, Ʃt 2g 2 (iii) Square d 2 x -y, dz 2, d xy, dyz b 1g, 1g, b 2g, Ʃb 1, Ʃa y planar and dxz e g Ʃb 2 Ʃe g 82

83 (b) (i) (ii) (iii) 3.(a) (i) (ii) E g * - t 2g t 2 * - E g B 1g * - A 1g ML -bonding It increase Crystal Field Splitting Energy E very much. (iii) [Fe(CN) 6 ] 4- (b) (i) Names (ii) Symmetry (iii) LGO dxy 1 2 x1 y2 x3 y4 dxz t 2g 1 2 x5 y1 x3 x6 dyz 1 2 x2 y5 x6 y 4 83

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 Correlation 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 Optical Rotatory Dispersion, ORD Circular Dichroism, CD. 3.9 Let us sum up 3.10 Check Your Progress : The Key 84

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 : 85

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 86

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 ( cm -1, 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 t 2g level to e g * levels belonging largely to the metal itself. (2) Ligand-to-Metal Charge Transfer Bands When the electron transition takes place from a MO located primarily on the ligand (M L bonding or orbitals) to a nonbonding or antibonding MO located primarily on the metal atom, the lgand-to-metal charge transfer bands are observed. These cannot be explained by CFT and represent the tendency of ligands to reduce the metal ion. The semiempiral MOT is adequate for explaining it. (3) Metal-to-Ligand Charge Transfer Bands These involve the transition of electron from an antibonding or non-bonding orbital, concentrated on the metal atom, to the antibonding orbital located primarily on the ligand, and measures the tendency of the metal ion to reduce the ligand. These bands are 87

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 d n is given in Table 3.2. In the absence of the external field, in fact the five d orbitals remain undefined. Under the influence of an octahedral field, these split into a group of t 2g and e g (or e g *) orbitals. An s orbital, being spherically symmetrical and nondegenerate, is not split into any states at all, and the p orbitals are not split in octahedral field as they interact with the field to the same extent. f orbitals get split into three levels a triply degenerate t 1 level, a triply degenerate t 2 level and a singlet level a 2 (Fig. 3.1). 88

89 3.3.1 Correlation Diagrams To develop energy level diagrams of complexes, start is made with the free ion of d n 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 d 1 configuration becomes 2D which breaks into two energy levels, 2 T 2 (corresponding to t 2g ) and 2 E g (corresponding to the e g ) levels. A d 2 configuration has two Russell Saunders states : a low energy 3F and the high energy 3 P state. These states behave, in an octahedral field, exactly in the same way as the f and p orbitals. Hence, a 3F state gives 3 T 1g, 3 T 2g and 3 A 2g states whereas the 3 P state is not split and gives 3 T 1g (P) state (Fig. 3.2). In d n configuration electron electron interaction starts. 89

90 Table 3.1 Splitting of the Terms in a d n ion in Weak Octahedral Field Configuration of free ion d 1 Ground state of free ion 2 D Energy level diagram A Configuration of free ion d 6 Ground state of the ion 5 D Energy level diagram A d 2 3 F B d 7 4 F B d 3 d 4 d 5 4 F 5 D 6 S inverted B inverted A no splitting d 8 d 9 d 10 3 F 2 D 1 S inverted B inverted A no spliting * With reference to Fig. 3.2 Note that the ground states of a d n 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 By considering the hole formalism, all the splitting, can be explained for a d n system in terms of Fig. 3.2 as A, B, inverted A or inverted B (Table 3.2, Fig. 3.3). Table 3.2 Splitting of the Ground State Terms in an Octahedral Field Term S P D F G H I A 1g T 1g Components in octahedral field E g + T 2g A 2g + T 2g + T 1g A 1g + E g + T 1g + T 2g E g + T g + T 1g + T 2g A 1g + A 2g + E g + T 1g + T 2g + T 2g 90

91 Fig. 3.3 Ground term splitting and energy (in Dq) unit for d n 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 d n 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 d n Ion Configuration d 1, d 2 d 2, d 8 d 3, d 7 d 4, d 6 d 5 Ground State Term 2 D 3 F 4 F 5 D 6 S Higher Energy Terms 3 P, 1 G, 1 D, 1 S 4 P, 2 H, 2 G, 3 F, 2 x 2 D, 2 P 3 H, 3 G, 2 x 3 F, 3 D, 2 x 3 P, 1 I, 2 x 1 G, 1 F, 2 x 1 D, 2 x 1 S 4 G, 4 F, 4 D, 4 P, 2 I, 2 H, 2 x 2 G, 2 x 2 F, 3 x 2 D, 2 P, 2 S 91

92 the energy terms (ground state as well as the excited state) should be shown (Fig. 3.4 for a d 2 ion) and analyzed. Fig. 3.4 Correlation diagram for d 2 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 Selection Rules There are two selection rules for electronic transitions in complexes; these are similar to the selection rules in atomic spectroscopy. They are : (a) Spin Multiplicity Rule : Transitions between states of different multicity, S, are forbidden. Usually this means that the number 92

93 of unpaired electrons must not be changed, but this is not quite the same thing. For example, the transition S 2 S 1 P 1 is spin-allowed so long as the spins of the two electrons in the S 1 P 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 d 5 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 d 5 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 6 S. Laporte Selection Rule : Transitions within a given set of p or d orbitals (i.e. transitions involving only redistribution of electrons in the given sub-shell) are forbidden if the molecule or ion has a centre of symmetry. A more formal statement of this rule, first put forward by Laporte, is that in a molecule which has a centre of symmetry, transitions between two g states or between two u states are forbidden (g = symmetrical; u = unsymmetrical). 93

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 [MX 4 ] 2- where M = Ni, Co, Cu, and X = Cl, Br, I is mm -1 mol -1 whereas in the octahedral complexes, these Laporte forbidden transitions have absorption of ca 5 units only. The vibrational and electronic coupling, called the vibronic coupling removes the centre of symmetry. If a forbidden band lies near an allowed band (due to permitted transition), the mixing of the fully allowed and forbidden energy states due to the vibronic coupling gives intense absorptions. This is called intensity stealing and depends upon the energy differences between the two states. 94

95 Check your Progress-1 Notes:(i) (ii) (a) Write your answers in the space given below. Compare your answers with those given at the end of the unit. The ground state Terms for the following free ions and their splitting components in an octahedral field are : d n ion Ground Term Splitting Components d d d (b) With what the two selection rules are related? Name the complex in which the rule is violated. Rule Related with Breakdown 1 st Rule nd Rule ORGEL DIAGRAMS Orgel diagrams are popular methods to represent the ground - and the excited states of a configuration. Similar to the correlation diagrams, in orgal diagrams also energies of states are plotted against the field strengths. Although Orgel diagrams are very simple, but as in these diagrams excited states of different multiplicity, other than the ground state, are not expressed, these diagrams are applicable to weak field cases only. 95

96 If one plots the magnitude of splitting of the energy levels with the increasing ligand field for a d n 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 d n ions in Fig. 3.5 Fig. 3.5 Orgel diagram for the P and F states in octahedral as well as in the tetrahedral field. Further, the energy level diagram for the d n ion in the tetrahedral field has the same form as that for the d 10-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 Co 2+ (d 7 ) in the tetrahedral and octahedral fields. Here again we can see the inverse relationship between these two symmetries. This is because the tetrahedral field, in fact is a negative octahedral field. In this diagram effects of the mixing of the terms can be seen. It is a common rule, only the terms of the same symmetry are mixed; while the limit of mixing is inversely proportional to the 96

97 difference in the energy between these terms. For Co 2+ these terms are of two, 4T 1 (tetrahedral) and 4T 1g (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 CO 2+ ) this is represented for the pairs of 4T 1g and 4T 1 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 4T 1 terms gradually closer with the increasing field strength. While in octahedral complexes just reverse of this is seen. Thus, in tetrahedral complexes the limit of mixing is high. Orgel diagrams are simple means to assess the number of spin allowed absorption bands for a complex in an ultraviolet or visible spectrum. 3.5 TANABE SUGANO DIAGRAMS In the strong field cases, the term values for the free ions fail to give the field splitting. Rather, the attempts are made to estimate the Racah parameters by placing electrons one by one in the t 2 and e g 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 d 4, 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). 97

98 Dq/B (a) Dq/B (b) Fig. 3.6 Tanabe-Sugano diagrams for (a) d 2 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. 98

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. [CoCl 4 ] 2- is blue whereas [Co(NH 3 ) 6 ] 2+ pink. Multiplicity forbidden transition bands are sharp, whereas the multiplicity allowed t 2g e g 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 d 6 ion in an octahedral stereochemistry. In these diagrams the ground state is taken as the abscissa and other energy states are recorded relative to this abscissa. While the interelectronic repulsion is expressed in terms of Racah parameters, B and C. The constant B is sufficient to give information about the energy differences in the levels of same spinmultiplicity, but for the terms of different multiplicity, both B and C parameters become necessary. For the energy difference between the 99

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 d 2, 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 d 5 ion is 5 D, which splits into high energy 5 E g and low energy 5 T 2g levels. The low spin state for d 6 ion is 1 I, which is a high energy term in the ground state. But as can be seen from the energy level splitting diagram for the d 6 ion, the 1 A 1g and 1 T 1g states get stabilized more than the 5 T 2g state so that at higher field strength, the ground state of the ion becomes 1 A 1g. The spectra of the cobalt (III), a d 6 ion can now be predicted easily. The high spin complex [CoF 6 ] 3- shows only one absorption band at cm -1 corresponding to 5 T 2g 6 E g transition, whereas the low spin cobalt (III) complexes show two peaks 100

101 corresponding to 1 A 1g 1 T 1g and 1 A 1g 1 T 2g as shown by both [Co(en) 3 ] 3+ (21400, cm -1 ) and [Co(Ox) 3 ] 3- (16500, 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 (B o ), the value of electronic-repulsion constant in its complex (B) is always less. This is due to nephelauxetic effect. The ratio of B and B 0 is known as Nephelauxetic constant, and is the measure of covalent character in the complex: B B o Valueof Bin the Complex Valueof Bin the freion Lesser is the value of, higher is covalent character in the complex. For example, the low spin complex of Co(III) (d 6 ion), [Co(en) 3 ] 3+, gives two bands at cm -1 and cm -1 in its spectrum (Fig. 3.8); which may be assign to: cm -1 = 1 A 1g 1 T 1g transition; and cm -1 = 1 A 1g 1 T 2g transition. The ratio of these energies will be:- 1 1 A 1g A 1g 1 T 1 T 2g 1g 29600cm 21550cm

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, 1 1 A1 g T1 g B cm 38 B This gives, B = 570 Cm -1, which very much less, than the value of Bo, for the free Co 3+ ion (= 1100 cm -1 ). From the value of / B (= 40) and B(=570 cm -1 ), the value of obtained is cm -1 (Theoretical value, cm -1 ). 102

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 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. complexes: Two types of the charge transfer (CT) bands appear in the metal (1) The LM CT bands as shown by oxide, chloride bromide and iodide complexes. For d 0 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 10 4 cm -1 in MnO4 - ion, assigned to the t 1 e *, t 2 e * and t 1 t 2 *, * transitions, give the 10Dq for the oxide ion as e * t 2 * as 2.59 x10 4 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 ( mm -1 mol -1 ). 103

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 t 2g orbitals to the antibonding orbitals of the diimine group. In the case of the 1, 10-phenanthroline complex Fe(phen) 2+ 3, 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 t 2g * (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 RuCl 2-6 and IrBr 2-6 (d 4 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 t 2g * and e g * orbitals of the metal atom. In IrBr 3-6 (a d 6 complex), the t 2g * orbitals are field, and only the transitions to the e g * orbitals can be observed. Similarly, ammine and halo (except fluoro) complexes of Co(II) and Cr(III) give charge-transfer bands at 250 mm or at the higher wave 104

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 ) 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 Optically Rotary Dispersion (ORD): The effect of a given concentration of a particular optically active species on the plane of polarization of light is not constant, but depends upon the wavelength of the light, the direction as well as the 105

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 n 1 and n r, for the left and right circularly polarized light. Hence, the two components interact differently with the medium and emerge out of phase. On combination, the plane of polarized light gets rotated. The variation of the rotation of the plane of plane polarized light with wavelength (Fig. 3.9) is called the optical rotatory dispersion, (ORD) and the abrupt reversal of the rotation in the vicinity of the absorption band is called the Cotton effect. Fig. 3.9 (a) Positive Cotton effect. (b) Negative Cotton effect. ORD curves are useful in the assignment of absolute configuration. For example, the configurations of the enantiomers of tris (ethylenediamine) cobalt (III), and bis (ethylenediamine) glutamatocobalt (III) are known from X-ray investigations, and it is found that the three A-(D)-configuration of the ORD spectra to one of known configuration. For the ORD spectrum to be unambiguous, no other absorption must be nearby. 106

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 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 (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 (D 3 ) symmetry. That difference, however, shows up in the CD spectra where an additional band is seen near cm -1. Fig (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 1 A 1 1 E and 1 A 1 1 A 2. The circular dichroism is the difference between the molar absorptivities of the left and right polarized 1 r, (solid curves in Fig. 3.9). Complexes having the same sign of CD for a given absorption band will have the same absolute configuration. Hence, CD spectra can be used 107

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(NH 3 ) 4 (en)] 3+, in practice it proves to be impossible to isolate them because of the rapid interconversion of the ring conformers. If two or more rings are present it one complex, they can interact with each other and certain conformations might be expected to be stabilized as a result of possible reductions in interatomic repulsions. For example, consider a square planar complex containing, two chelated rings of ethylenediamine. From a purely statistical point of view we might expect to find three structures, which may be formulated M, M and M (which is identical to M ). The first two molecules lack a plane of symmetry, but M is a meso form. Corey and Bailar were the first to show that the M and M should predominate over the meso form. Check Your Progress- 3 Notes :(i) Write your answers in the space given below. (ii) Compare your answers with those given at the end of the unit. 108

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: 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 t 2g level to e* g levels largely in the metal itself. 109

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 d n configuration are: d 1 2 D d 6 5 D d 2 3 F d 7 4 F d 3 4 F d 8 3 F d 4 5 D d 9 2 D d 5 6 S d 10 1 S In presence of the external field the degenerate d-orbitals split into groups of orbitals according to the field strengths. Orgel and Tanabe-Sugano diagrams represent the ground and excited states of a given configuration. Orgel diagram plot the magnitude of splitting of the energy levels with the increasing ligand for d n system, and take into consideration the spin orbit coupling and mixing of the different energy states. It plots the energy, E, against the field strength, B. Orgel diagrams fail to explain the spectra fully, especially in cases of strong fields. Tanabe-Sugano diagrams are most suitable for explaining spectra both in cases of weak and strong field. They plot the energies in 110

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, B o : B B o From the Tanabe-Sugano diagrams, the value of E/B for the lowest energy transition may be read (=x). This will give the value of B: Lowest energy B transition X From the values of /B (=X) and B, the value of is obtained. There are two phenomena associated with the d-d transition that are useful in assigning absolute configuration in optically active metal chelates, Optical Rotatory Dispersion (ORD) and Circular Dichroism (CD). The variation of the rotation of the plane of plane polarised light with wave length is called ORD. While CD spectrum is a plot of the difference of the molar absorption coefficients for right and left circularly polarized light against wave length. ORD and CD form the basis for cotton effect (The abrupt reversal of the rotation in the vicinity of the absorption band), and as a rule if in analogous compounds, corresponding electronic transitions 111

112 show cotton effects of same sign the compounds have the same optical configuration CHECK YOUR PROGRESS : THE KEY 1(a) d n ion Ground Term Splitting Components d 1 2 d E g and T 2g d 3 4 F A 2g, T 2g and T 1g d 5 6 S A 1g (No Splitting) (b) 1 st rule Spin multiplicity Spin free oh complexes of Mn 2+( d 5 ) 2 nd rule Redistribution of NiCl 2-4, Th complex gives more electrons in a intense absorption than of an Oh given sub-shell complex. 2(a)(i) In Orgel diagrams energy state E are plotted against field strength. B; while in Tanabe-Sugano diagrams E/B is plotted against /B. (ii) Orgel diagrams are suitable for complexes of weak fields only; but tanabe-sugano diagrams are suitable for both weak and strong fields. (b) B and C Racah parameters of complexes represent inter electronic repulsion; while nephelauxetic coefficient represent nephelauxetic effect and is the ratio of B and B 0 i.e. B B o 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. 112

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 t 2g orbitals to anti-bonding orbitals of the ligand. (b)(i) ORD is the variation in sign and magnitude of rotation of plane of polarized light with wave length; while CD plots the difference of the molar absorption coefficients for right and left circularly polarized light against wave length. (ii) The absolute configuration in optically active metal chelates may be assigned using a general rule: if in analogous compounds, corresponding electronic transition shows cotton effects of same sign, the compounds have the same optical configuration. 113

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 114

115 UNIT-4 MAGNETIC PROPERTIES OF TRANSITION METAL COMPLEXE Structure 4.0 Introduction 4.1 Objectives 4.2 Magnetic Moments Number of Unpaired Electrons Spin Only Formula 4.3 Anomalous Magnetic Moments Orbital Contribution in Magnetic Moments Curie's Law 4.4 Magnetic Exchange Coupling 4.5 Let Us Sum Up 4.6 Check Your Progress: The Key 115

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. 116

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", x m K x d where d and M are the density and molecular weight of the substance, respectively. For most substances, K, x and x M 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 x M have relatively large x M K M d 117

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; where x corr M 2 N 3kT corr X M is the molar susceptibility which has been corrected both for the diamagnetic contribution to the susceptibility (due to the non paramagnetic atoms in the sample) and for any small temperatureindependent paramagnetism arising from paramagnetic excited states of the system. N is Avogadro's number, k is the Boltzmann constant, µ is the "magnetic moment" of the molecule, and T is the absolute temperature. By substituting numerical values for N and k, we obtain; X corr M T or 2.83 X corr T M Number Of Unpaired Electrons Once an experimental value of x M has been obtained for a paramagnetic substance, it can be used to determine how many unpaired electrons there are per-molecule or ion. In order to translate the experimental result into the number of unpaired spins, it must be recognized that a measured susceptibility will include contributions from both paramagnetism and diamagnetism in the sample. Even though the latter will be small, it is not always valid to consider it negligible. The most common procedure is to correct a measured susceptibility for the diamagnetic contribution. Compilations of data from susceptibility 118

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: N x M 3RT 2 2 where N is Avogadro's number, R is the ideal gas constant, T is the absolute temperature, and µ, is expressed in Bhor magnetrons (BM) (1 BM = eh/4 m). Solving this expression for the magnetic moment gives: 3RTX N 1/ 2 m 1/ ( x T) 2 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, orbitalorbital, and spin-orbital. For some complexes, particularly those of the lanthanides, we must consider all three types of coupling. The theoretical paramagnetic moment for such a complex is given by: g[ J( J 1)] 1/ 2 119

120 where J is the total angular momentum quantum number and g is the Lande splitting factor for the electron, defined as: g 1 J( J 1) S( S 1) L( L 1) 2J( J 1) The value of J depends on the total orbital angular momentum quantum number, L, and total spin angular momentum quantum number Spin Only Formula For complexes in which spin-orbit coupling is nonexistent or negligible but spin and orbital contributions are both significant, the predicted expression for µ is; [ 4S( S 1) L( L 1)] 1/ 2 This equation describes a condition that is never fully realized in complexes because the actual orbital contribution is always somewhat less than the ideal value. This occurs because the orbital angular momentum is reduced from what it would be in the free metal ion by presence of ligands. In the extreme case, where general situation in complexes having A or E ground states, which would include octahedral d 3, 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; 1/ 2 [ 4S( S 1) 2[ S( S 1)] 1/ 2 which is known as the spin only formula for magnetic moment. By recognizing that S will be related to the number of unpaired electrons (n) by S = n/2, the expression may be further simplified to; [ n( n 2)] 1/ 2 120

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 x M, is given by the relation; x M = Magnetic moment, µ is given by the relation; µ = The spin only formula is; µ = Magnetic moment, µ and number of un-paired electrons, n, are related as; µ = ANOMALOUS MAGNETIC MOMENTS Table 4.1 indicates that the values of magnetic moment calculated using the spin only formula in number of cases differ from the values obtained from theoretical considerations. This difference is supposed to be due to two reasons, firstly due to the contribution of orbital magnetic moment; and secondly due to dependence of magnetic properties on the experimental temperature (Curie's Law). 121

122 Central metal Table 4.1: Magnetic properties of some complexes of the first-row transition metals No. of d electrons High spin complexes No. of unpaired electrons µ(expt) BM µ(calc) b BM Low spin complexes No. of unpaired electrons µ(expt) BM µ(calc) b BM Ti V V V Cr Mn Cr Mn Mn Fe Fe Co Co Ni Ni Cu

123 4.3.1 ORBITAL CONTRIBUTION IN MAGNETIC MOMENTS In an octahedral ligand field, only the t 2g orbitals remain degenerate and rotationally related. The e g orbitals get separated by 10Dq. 2 Hence, orbital momentum due to the d 2 x -y 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 d 1, d 2, d 6 and d 7 ion complexes and low spin d 4, d 5 ions in octahedral field. In the tetrahedral field, the high spin d 3, d 4, d 8 and d 9 ion should have a significant contribution from the orbital angular momentum. The magnetic moment of [CoCl 4 ] 2- (4.4 BM) and that of [Co(H 2 O) 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 µ eff values obtained for Cr 2+, Cr 3+, V 3+ and V + and higher values for spin free Fe 2+, Co 2+, Ni 2+ and Cu 2+ complexes. Greatest deviations occur for the Co 2+ and Fe 2+ complexes, for which unquenched orbital moments contribute significantly. For the 4d and 5d ions, diamagnetism results for even numbered electrons, and paramagnetism to the extent of one unpaired electron only is observed for the old numbered electrons, indicating that spin pairing takes place for these ions as far as possible. This may be due to (i) reduced enterelectronic repulsions in larger sized ions reducing the 123

124 electron pairing energies, (ii) higher LF or MO splittings. The µ at room temperature is generally lower than µ s 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 MnO - 2-4, CrO 4 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 CURIE'S LAW The observed magnetic moments for the metals in t 2g ground state are temperature dependent and usually depart from the µ s value due to probably the t 2g electron delocalization and lower symmetry ligand field components. Pierre Curie established in 1895 that paramagnetic susceptibility is inversely proportional to the absolute temperature. x M = C/T This expression, which is known as Curie's Law, is actually a restatement of magnetic moment. The Curie law is obeyed fairly well by paramagnetic substances that are magnetically dilute, i.e. those in which 124

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; x M C ( T ) where ø is a constant with units of temperature. If the interacting magnetic dipoles on neighboring atoms tend to assume a parallel alignment, the substance is said to be ferromagnetic (Fig. 4.1(b)). If, on the other hand, the tendency is for an anti-parallel arrangement of the coupled spins, the substance is anti-ferromagnetic.(fig. 4.1(c)) In any material that exhibits magnetic exchange, the tendency towards spin alignment will complete with the thermal tendency favoring spin randomness. In all cases, there will be same temperature below which magnetic exchange dominates, this temperature is called the Curie temperature (T C ) if the type of exchange displayed is ferromagnetic and the Neel temperature (T N ) if it is anti-ferromagnetic. The change in susceptibility as the temperature is decreased below either T C or T N may be quite dramatic. Paramagnetism Ferromagnetism Antiferromagnetism (a) (b) (c) Fig.4.1 Schematic representations of magnetic dipole arrangements in (a) paramagnetic, (b) ferromagnetic, and (c) antiferromagnetic materials. Fe(Phen) 2 (CNS) 2 is an example which shows significant variation in magnetic moment with temperature. (Fig. 4.2) 125

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 energy difference between. For complexes in which the 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 Fe 2+ ion. At the two extremes, it forms high spin paramagnetic [Fe(H 2 O)] 2+ diamagnetic [Fe(CN) 6 ] 4- (S=0). (S=2) and low spin Octahedral complexes with 4, 5, 6 or 7 d electrons can be either high-spin or low-spin, depending on the magnitude of the ligand-field splitting,. When the ligand-field splitting has an intermediate value such that the two states of the complex have similar energies, the two states can coexist in measurable amounts at equilibrium. Many "crossover" systems of this type have been studied. 126

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 d 6 complexes show that near the crossover point between weak and strong field the difference in energy between the spinfree ( 5 T 2g ) and spin-paired ( 1 A 1g ) ground states becomes very small (Fig. 4.3) within this region, it is reasonable to expect that both spin state may be present simultaneously and that the degree to which each is represented will depend on the temperature ( - P = kt). A complex illustrating these effects is [Fe(phen) 2 (NCS) 2 ] (Fig. 4.2). At high temperature a moment consistent with four unpaired electrons is observed, but as the temperature is decreased, a sharp drop in magnitude is observed at 175K where the low-spin form becomes dominant. Usually spin transitions occur somewhat more gradually than in the case shown here, and reasons for the abruptness observed for this complex, as well as some residual paranagnetism seen at low temperature have been discussed extensively. 5 T 2g E 1 A 1g Fig.4.3 Variation in energies of 5 T 2g and 1 A 1g terms with increasing for d 6 octahedral complexes. At weak field (high spin complexes) the ground term is 5 T 2g, while at strong fields (low spin complexes) it is 1 A 1g Note that in the region immediately on each side of the spin crossover point, the energy difference between the two terms is small; thus high and low spin complexes coexist. 127

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 d n configurations (n=...) the two states (low-spin and high-spin) of complexes can coexist in measurable amount's at equilibrium at ligand field splitting has. (D) The change from high spin to low spin can be explained in terms of... (E) The crossover point is reached when the difference in energy between...states become very small. 128

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, X M as- X K d or X M K M d where d and M are the density and molecular weight of the substance, respectively. x M is the molar susceptibility which has been corrected both for the diamagnetic contribution to the susceptibility and for any smll temperature-independent paramagnetism from paramagnetism excited states of the system, and may be given as, X corr M T or 2.83 X corr T M 129

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 X M has been obtained for a paramagnetic substance, it can be used to determine how many unpaired electrons there are per molecule or ion. For complexes in which spin-orbit coupling is nonexistent or negligible but spin and orbital contributions are both significant µ is given by; μ [4S(S 1) L(L 1)] 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; 1/ 2 μ [4S(S 1) 1/ 2 2[S(S 1)] 1/ 2 By recognizing that S will be related to the number of unpaired electrons (n) by S = n/2, the above expression is simplified to; μ [n(n 2)] 1/ 2 Magnetic moment calculated using the spin only formula in number of cases differ from the values obtained from theoretical considerations. This deviation may due to be either the contribution of orbital magnetic moment, or due to dependence of magnetic properties on the experimental temperature, (Curie's Law). The orbital angular momentum formula may be important for the high spin d 1, d 2, d 6 and d 7 ion complexes, and low spin d 4, d 5 ions in octahedral field. 130

131 The observed magnetic moments for the metals in t 2g ground state are temperature dependent and usually depart from the µ s values due to probably the t 2g electron delocalization and lower symmetry ligand field components. The Curie's Law states that paramagnetic susceptibility is inversely proportional to the absolute temperature; x M = C/T If the interacting magnetic dipoles on neighboring atoms tend to assume a parallel alignment, the substance is said to ferromagnetic, and if, on the other hand, the tendency is for an anti-parallel arrangement of the coupled spins, the substance is antiferromagnetic. Fe(phen) 2 (CNS) 2 is an example which shows significant variation in magnetic moment with temperature. A number of transition metal ions form both high and low spin complexes. These two spin configurations in octahedral complexes are explained in terms of relative magnitudes of and pairing energy (P). High spin complexes are formed when < P and low spin when > P. For complexes in which the energy difference between and P is relatively small an intermediate field situation, it is possible for the two spin sates to co-exist in equilibrium with each other. Variation in energies of 5 T 2g and 1 A 1g terms with increasing for d 6 octahedral complexes show, at weak fields (high spin complexes), the ground term is 5 T 2g, while at strong fields (low spin complexes) on each side of the spin crossover point, the energy difference between the two terms is small; thus high and low complexes may coexist. 131

132 4.7 CHECK YOUR PROGRESS: THE KEY 1. (i) KM X M d (ii) = 2.84 (X M T) ½ (iii) = [4S (S+1) ] ½ = 2 (S(S+1)] ½ (iv) = [n (n+2)] ½ 2.A (i) High spin d 1, d 2, d 6 and d 7 ion complexes, and low spin d 4 and d 5 ions. (ii) For Fe 2+ and Co 2+ complexes. B. Curie's Law states that paramagnetic susceptibility is inversely proportional to the absolute temperature, i.e. X M = C/T C. For d n 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 5 T 2g and spin paired 1 A 1g ground states become very small. 132

133 UNIT-5 METAL π COMPLEXES Structure 5.0 Introduction 5.1 Objectives 5.2 Metal Carbonyls Classification Isolobal Concept Methods of Preparation and Properties Structure Vibrational Spectra 5.3 Metal Nitrosyls Neutral NO and NO - Complexes Complexes of NO Pure Nitrosyl Complexes Nitrosyl Carbonyl Complexes Nitrosyl Halide Complexes Nitroso Cyanide Complexes 5.4 Dinitrogen Complexes Fixation of Nitrogen 5.5 Dioxygen Complexes Heme Proteins and Transportation of O Haemoglobin 5.6 Tertiary Phosphine as Ligand 5.7 Let Us Sum Up 5.8 Check Your Progress: The Key 133

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: 134

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 examplenickel forms volatile nonpolar carbonyls, where as alkali and alkaline earth metals from non-volatile electrovalent carbonyls. 135

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 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. [Fe 3 (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 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; Ni 40 C 4CO Ni( CO ) 4 136

137 Electrons needed to attain noble gas configuration First transition series Second transition series Third transition series 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 V(CO) 6 blue solid Cr(CO) 6 white solid (sublimes) Mo(CO) 6 white solid(subli mes) W(CO) 6 white solid(subli mes) Mn 2 (CO) 10 yellow solid (m.p.154 o ) Te 2 (CO) 10 white solid Re 2 (CO) 10 white (m.p.177 o ) solid Fe(CO) 5 yellow liquid (b.p.103 o ) Fe 2 (CO) 12 black solid (sublimes) Ru(CO) 5 colourless liquid 22 o ) Ru 2 (CO) 9 * Ru 3 (CO) 12 orange (m.p.-154 o ) Os(CO) 5 colourless liquid (m.p.15 o ) Os 2 (CO) 9 Orange Os 2 (CO) 12 Yellow (m.p.224 o ) (m.p.- solid CO 2 (CO) 8 orange solid (m.p.51 o ) CO 6 (CO) 16 black solid Rh 2 (CO) 8 * Rh 4 (CO) 12 orange solid Rh 6 (CO) 16 black solid solid solid solid Ir 2 (CO) 8 yellow solid Ir 4 (CO) 12 yellow solid Ir 6 (CO) 16 red Ni(CO) 4 colourless liquid (b.p.43 o ) 137

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: C 6 H 5 MgBr + CrCl 3 + CO Cr(CO) 2 (C 6 H 5 ) + MgBrCl + MgBr 2 The unstable intermediate compound is composed with acid to yield the hexacarbonyl: 3Cr(CO) 2 (C 6 H 5 ) 4 + 6H Cr(CO) 6 + 2Cr C 6 H H 2 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. 2NiX 4 + 2nCO 2Ni(CO) n X + X 2 Ni(CO) n X + (4-2n)CO Ni(CO) 4 + NiX 2 ii.ruthenium pentacarbonyl may be prepared by the action of CO and Rul 3 in the presence of an iodine acceptor: RuI [ CO] [ CO] 3 Ru( CO) 4 I2 Ru( CO) 5 Similarly [Ir(CO) 4 ] 7 may be prepared. d. Synthesis by carbonylating the metallic salts with CO in the presence of reducing agent. When salts like R 4 I 3, CrCl 3, VCl 3 are made to treat with CO in presence of a suitable reducing agent like Mg, Ag, Cu, Na, H 2 etc. 138

139 o CrCl3 CO LiAlH y 2Rul CO + 6Ag 2Mnl CO + 2Mg 115 Cr(CO) 6 + LiCl + AlCl C250atmpressur e 2Ru(CO) 5 + 6Agl 25 C210atmpressur e 2Mn 2 (CO) Mgl e. Synthesis from other carbonyls: when iron pentacarbonyl is exposed to UV light it loses CO and forms Fe 2 (CO) 9. This compound undergoes thermal decomposition to yield iron pentacarbonyl and trimeric tetracarbonyl. 2Fe(CO) 5 2Fe 2 (CO) 9 U.V Fe2(CO) 4 + CO. heat Fe(CO) 5 + [Fe(CO) 4 ] 3 + CO f. Synthesis from Carbonyl hydrides: when iron carbonyl hydride is oxidised by MnO 2 or H 2 O 2, [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. OsO 4 + 9CO 100 C Os(CO) + 4CO atm pressure Re 2 O CO 75 C Re 2 (CO) CO atm h. Preparation of Mo (CO) 6 and W(CO) 6 from Fe(CO) 5 MoCl 6 + 3Fe(CO) 5 Mo(CO) 6 + 3FeCl 2 + 9CO i. Preparation of Fe 2 (CO) 9 and Os 2 (CO) 9 from Fe(CO) 5 and Os(CO) 5-9 with cooled solution of Fe(CO) 5 and Os(CO) 5 in glacial CH 3 COOH is irradiated with u.v. light, Fe(CO) 9 and Os 2 (CO) 9 are obtained respectively. 2Fe(CO) 5 2Os(CO) 5 U.V. light U V. light Fe 2 (CO) 9 + CO. Os 2 (CO) 9 + CO 139

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 CCl 4. 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, ( o C) V(CO) 6 Black crystals Decomposes at 70 o C, 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 Mn 2 (CO) 10 Golden crystals 154 o -155 o Re 2 (CO) 10 Colourless crystals Sublime at 140 o and decompose at 177 o C Fe(CO) 5 Yellow Liquid B.P. 103 o C Fe 2 (CO) 9 Bronze Mica-like Decomposes at 100 o C platelets Fe 3 (CO) 12 Dark green crystals Decomposes at 140 o C CO 2 (CO) 8 Orange crystals 51 o C Ni(CO) 4 Colourless Liquid B.P. 43 o C 140

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) PR 3, PCl 3, Py, CH 3 OH etc. Ni(CO) 4 + 4CNR Ni(CNR) 4 + 4CO Ni(CO) 4 + 4PCl 3 Ni(PCl 3 ) 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 ] - Na 2 Co 3 + H 2 O 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 NH 3 is able to convert Fe 2 (CO) 9., Co 2 (CO) 8 Fe 3 (CO) 12, Cr (CO) 6,. Mn 2 (CO) 10 etc, into carbonylate anions and in this conversion these carbonyls are reduced. Fe 2 (CO) 9 + 4Na 2Na 2 [Fe 2- (CO) 4 ] 2- + CO Co 2 (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 + X 2 Fe(CO) 4 X 2 + CO Mo(CO) 6 + Cl 2 Mo(CO) 4 Cl 2 + 2CO Mn 2 (CO) 10 + X 2 (X = Br, I) 2Mn(CO) 5 X Both Co 2 (CO) 6 and Ni(CO) 4 are decomposed into metallic halides and CO when treated with halogens. Co 2 (CO) 8 + 2X 2 2CoX 2 + 8CO Ni(CO) 4 + Br 2 NiBr 2 + 4CO 141

142 4. Action of NO: many carbonyls react with nitric oxide (NO) to form metal carbonyls nitrosyls. For example: Fe(CO) 5 + 2NO 95 C Fe(CO) 2 (NO) 2 + 3CO 3Fe 3 (CO) 9 +4NO 2Fe(CO) 2 (NO) 2 + Fe(CO) 5 +Fe 3 (CO) 12 +6CO 5. Action of H 2 : Formation of carbonyl hydrides (reduction): when Mn 2 (CO) 10 and Co 2 (CO) 8 react with H 2, they get reduced to carbonyl hydrides, Mn(CO) 5 H and Co(CO) 4 H respectively. Mn 2 (CO) 10 + H 2 Co 2 (CO) 8 + H 2 e C, 200atmpressur 200 2[Mn - (CO) 5 H + ] C, 200atmpressur e 2[Co - (CO) 4 H + ] 6. Action of heat: Different carbonyls yields different products when heated for example: Fe(CO) 5 3Fe 2 (CO) 9 Fe 3 (CO) 12 C 250 Fe + 5CO 70 C 3Fe(CO) 5 + Fe 3 (CO) 12 C 140 3Fe + 12CO Metal Carbonyls of Different Groups: 1. Carbonyls of Sixth B Metals These form carbonyls of one type only M(CO) 6 where M = Cr,Mo, Or W, but chromium also forms Cr(CO) 5+ A. Chromium Hexacarbonyl Cr(CO) 6. Preparation: i. It is prepared by job's method by passing CO at 50 atm. pressure and at room temperature into a suspension of chromic chloride in ether. Which has been treated with phenyl magnesium bromide at -70 o C. ii. Chromium hexacarbonyl can be prepared by treating a solution of a chromic salt dissolved in ether with Al(C 2 H 5 ) 3 and carbon mono-oxide at a high temperature and pressure. 142

143 iii. It may also be prepared by carbonylating CrCl 3 with CO in the presence of a reducing agent like LiAlH 4. CrCl 3 + CO + LiAlH 4 Properties: C, 200atmpressur e 175 Cr.(CO) 6 + LiCl + AlCl 3 1. Chromium hexacarbonyl exists in colourless rhombic crystals which sublime without decomposition and dissolve in either, chloroform, CCl 4 and benzene. 2. It is attacked by air, bromine, cold aqueous alkali, dilute acids conc. HCl and Conc.H 2 SO 4. It is however decomposed by chlorine or by conc. nitric acid. 3. Decompositions: It gets decomposed by F 2 at -75 o C to form CrF Action of Na-Metal or NaBH 4 : Cr(CO) 6 when is treated with Na metal or NaBH 4 in liq. NH 3 carbonylate anion is formed. In these reactions the carbonyls are reduced. Cr(CO) 6 + 2Na Cr(CO) Liq.NH NaBH4 / Liq. NH Na 2 [Cr 2- (CO) 5 ] 2- + CO Na 2 [Cr 2- (CO) 10 ] CO 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 py Cr 2 (CO) 7 (Py) 5 Cr( CO) 3 ( Py) 3 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 MoCl 6 or WCl 6 with CO in the presence of phenyl magnesium bromide. 143

144 2. Both may also be prepared by the action of CO at 225 o and 200 atm. pressure on metallic molybdenum or tungsten reduced in the presence of copper or iron. Properties: 1. They are colourless, Mo(CO) 6 sublimes at 40 o C and boils at o, whereas W(CO) 6 sublimes at 50 o C and boils at 175 o C. 2. They are stable in air and dissolve in organic solvents like ether, chloroform, CCl 4 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) 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) 5 Pyr 2 M 2 (CO) 7 PYr 5 M(CO) 3 PYr 3 2. Carbonyls or VII Group: These form volatile carbonyls of the formula M 2 (CO) 12 where M = Mn, Te and Re. Manganese carbonyl, Mn 2 (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. 2MnCl 2 +10CO+2Mg (in diethyl ehter) e 250 C,210atmpressur Mn 2 (CO) MgI 2 2. By carbonylating MnCl 2 with CO in presence of (C 8 H 5 ) 2 CONa 2MnCl 2 +10CO+4(C 8 H 5 ) 2 CONa. 165 C140 atm Mn 2 (CO) (C 6 H 5 ) 2 CO+4NaCl 144

145 Properties: 1. Manganese carbonyl forms volatile, golden yellow, crystalline, solid which melts at 155 o C in a sealed tube. It is soluble in organic solvents. It is slowly oxidised in air, especially in solution. 2. Action of halogens: Mn 2 (CO) 10 reacts with halogens to form carbonyl halides. Mn 2 (CO) 10 + X 2 (X=Br 2 I) 2Mn(CO) 5 X 3. Action of Na-Metal: Na-metal in liquid NH 3 converts Mn 2 (CO) 10 into carbonylate anion. In this reaction the oxidation state of Mn decreases from zero to -1; Mn 2 (CO) 6 + 2Na Liq.NH 2Na [Mn - (CO) 5 ] Action of H 2 : Mn 2 (CO) 10 gives carbonyl hydride, Mn(CO) 5 H; in the formation of this compound the oxidation state of Mn decreases from zero to -1. Mn 2 (CO) 10 + H 2 e 200 C, 200atmpressur 2[Mn - (CO) 5 H + ] 0 5. Substitution Reaction: Mn 2 (CO) 10 reacts with PR 3 to form Mn(CO) 4 (PR 3 ): Mn 2 (CO) 10 + PR 3 2Mn(CO) 4 (PR 3 ) + 2CO 6. Diamagnetic nature: Mn 2 (CO) 10 is a diamagnetic substance, diamagnetic character is confirmed by the fact that all the electrons in Mn 2 (CO) 10 are paired and Mn-Mn bonds is present in it. 3. Carbonyls of VIII Group Metals A. Carbonyls of Iron: Three carbonyls of iron are known, these are: a. Iron Pentacarbonyl, Fe (CO) 5 145

146 Preparation: i. It can be prepared by the action of CO on iron powder at 200 o C 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. FeI 2 + 4CO C 200 Fe (CO) 5 I 2 iii. It may also be prepared by the action of CO on FeS at 200 o C and 200 atm. pressure in the presence of copper. 2FeS + 10CO + 2Cu e 200 C, 200atmpressur 2Fe(CO) 5 + Cu 2 S Properties: i. Fe(CO) 5 is a yellow liquid which is soluble in methyl alcohol, ether, acetone and C 6 H 6. It is insoluble in H 2 O. ii. Decomposition: M thermal decomposition at 250 o C it yields pure Fe. 2Fe(CO) C Fe + 5 Co iii. Action of u.v. light: When cooled solution of Fe(CO) 5 in glacial CH 3 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 H 2 O and acids Fe (CO) 5 + H 2 SO 4 FeSO 4 + 5CO + H 2 v. Action of alkali: Fe(CO) 5 + 4NaOH Na+[Fe 2- (CO) 4 H + ] - + Na 2 CO 3 + H 2 O vi. Action of NH 3 : with NH 3 it yields Fe(CO) 4 H 2 Fe(CO) 5 + H 2 O + NH 3 Fe (CO) 4 H 2 + NH 2 COOH 146

147 vii. Reaction with halogen: Fe(CO) 5 + X 2 Fe (CO) 4 X 2 + CO The velocities of these reactions have been found to follow the order CI < Br < I. b. Iron Enneacarbonyl, Fe 2 (CO) 9 When iron pentacarbonyl is dissolved in glacial acetic acid and is exposed to u.v. light for 6 hours, Fe 2 (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 Fe 2 (CO) 9 + CO Properties i. Fe 2 (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 50 o C decomposes to form Fe 2 (CO) 12. 3Fe 2 (CO) 9 3Fe (CO) 5 + Fe 3 (CO) 12 ii. Action of heat: when heating is done at 100 o C, Fe 2 (CO) 9 decomposes to form iron, CO and some Fe (CO) 12. 4Fe 2 (CO) 9 Fe + Fe(CO) 5 + Fe 3 (CO) 12 + CO iii. Action of NO: With NO it gives Fe(CO) 2 (NO) 2 together with Fe(CO) 5 and Fe 2 (CO) 12. 3Fe 2 (CO) 9 +4NO2Fe(CO) 2 (NO) 2 +Fe(CO) 5 + Fe 3 (CO) CO c. Iron Dodecarbonyl, Fe 3 (CO) 12. Preparation It can be prepared by heating Fe 2 (CO) 9 dissolved in toluene at 70 o C. 3Fe 2 (CO) 9 3Fe(CO) 5 + Fe 3 (CO)

148 Properties i. Fe 3 (CO) 12 forms deep crystals which are soluble in organic solvents like toluene, alcohol, ether and pyridine. ii. Action of Heat: When heated to 140 o C Fe 3 (CO) 12 decomposes to give metallic iron and CO. Fe 3 (CO) 12 C 140 3Fe + 12CO iii. Reaction with Na: Carbylate axion is formed when Fe 3 (CO) 12 reacts with Na metal in Liq. NH 3. Fe 3 (CO) Na 3.NH Liq 3Na 2 + [Fe 2- (CO) 4 ] 2 iv. Substitution Reactions: This reaction takes place with pyridine and methyl alcohol. Fe 3 (CO) Py Fe 3 (CO) 9 (Py) 3 + 3Fe(CO) 5 B. Carbonyls of Cobalt It forms two carbonyls i. Cobalt Octacarbonyl, CO 2 (CO) 8 Preparation i. It is prepared by the action of CO and the reduced metallic cobalt at 220 o C and 250 atm. 2Co + 8CO Co 2 (Co) 8 ii. When a solution of cobalt carbonyl hydride is treated by an acid, hydrogen is evolved and Co 2 (CO) 8 remains 2Co(CO) 4 H Co 2 (CO) 8 + H 2 Properties 1. Cobalt octacarbonyl forms orange transparent crystals. 2. It is insoluble in water but is soluble to some extent in alcohol, ether, CS, etc. 148

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. NH 3 : When CO 2 (CO) 8 reacts with Na-metal in liq. NH 3, it gets reduced to carbonylate anion. Co 2 (CO) 8 + 2Na 3 lig NH 2Na[Co(CO) 4 ] 5. Action of NO: Co 2 (CO) 8 reacts with NO at 40 o C to form cobalit carbonyl nitrosyl, [Co - (CO) 3 (NO)] 0. Thus in this reaction the oxidation state of cobalt decreases from 0 to -1. Co 2 (CO) 8 + 2X 2 2Co 2 X 2 + 8CO 6. Dispropotination Reaction a. Strong bases cause disproportination into Co(+2) and Co(-1) 2Co(CO) NH 3 2[Co(NH 3 ) 6 ][Co(CO) 4 ] 2 + 8CO b. With isocyanides it gives penta-co-ordinate cobalt(i) cation. Co 2 (CO) 8 + 5CNR [Co(CNR) 5 ][Co(CO) 4 ] +4 CO (ii) Dodecarbonyltetra Cobalt, [CO 4 (CO) 12 ] Preparation 1. It is prepared by heating Co 2 (CO) 8 at 60 o C. 2. It may also be obtained by oxidizing cobalt carbonyl hydride below -26 o C. Properties: i. It is black crystalline solid. ii. It is very unstable easily oxidized by air and can be recrystallized from hot benzene. 149

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 o C. Ni + 4CO Ni(CO) 4 ii. When Nickel iodide is heated with CO in the presence of a halogen acceptor, nickel carbonyl is formed. NiI 2 + 4CO Ni(CO) 4 + I 2 Properties i. It is colourless liquid, m.p. = -23 o C, b.p. = 43 o C ii. It has no solubility in water but dissolved in organic solvents. iii. It decomposes at 180 o -200 o C in to nickel and CO. Ni(CO) C Ni + 4CO iv. It reacts with H 2 SO 4 and form NiSO 4 Ni(CO) 4 + H 2 SO 4 NiSO 4 + H 2 + 4CO v. It reacts with Ba(OH) 2 and gives BaCO 3 Ni(CO) 4 + Ba(OH) 2 H 2 Ni(CO) 3 + BaCO 3 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 200 o C and 200 atm. pressures Ru + 5CO Ru(CO) 5 150

151 Properties i. It is colourless soluble liquid having m.p. = -22 o C ii. It has no solubility in water but is soluble in alcohol, benzene and CHCl 3. iii. It undergoes decomposition to give Ru 2 (CO) 9 and Ru 3 (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, Ru 2 (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) 2 I 2. c. Ru 3 (CO) 12 It is prepared in small quantities along with Ru 2 (CO) 9 when Ru(CO) 5 is heated at 50 o C 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. Os 2 (CO) 5 It is a colourless having m.p. -15oC. It is obtained. i. by the action of CO on OsI 3 at 120 o C and 200 atm. pressure in the presence of copper. ii. by the action of CO on OsO 4 at 100 o C and 50 atm. pressure. OsO 4 + 9CO Os(CO) 5 + 4CO 2 b. Osmium eneacarbonyl, Os 2 (CO) 9 It is a yellow crystalline solid. It is prepared by the reaction of OsI 3 with CO in the presence of copper, it is more stable towards heat than Ru 4 (CO) 9. It melts at 224 o C and sublimes without decomposition. 151

152 F. Carbonyl Of Iridium It forms 2 carbonyls: a. Iridium Octacarbonyl, Ir 2 (CO) 8 It is prepared by the reaction of either KIr 2 Br 6 or KIr 2 Br 6 or KIr 2 I 6 with CO at 200 o C and 200 atm. pressures. It is yellow crystalline solid having m.p. 160 o C. b. Iridium Dodecarbonyl, Ir 4 (CO) 12 It forms orange yellow rhombohedra crystals which decomposes at 200 o C. It is prepared by treating Irl 3 with CO under pressure. G. Carbonyl of Platinum Preparation i. When CO is passed over PtCl 2 at 250 o C, PtCl 2 (CO) and 2PtCl 2 3CO are obtained on heating these yield PtCl 2 (CO) 2. 3PtCl 2 + 5CO PtCl 2.2CO + 2PtCl 2.3C0 Properties These carbonyls are decomposed by water and HCl. PtCl 2.CO + H 2 O Pt + 2HCl + CO 2 PtCl 2.CO + H 2 O Pt + 2HCl + CO 2 + CO PtCl 2.CO + HCl H[PtCl 3.CO] PtCl 2.CO + HCl H[PtCl 3.CO] + CO Structure of Metal carbonyls 1. Effective Atomic Number Rule: The structure of CO is :C:O: It is probable that the lone pair of electrons on the carbon atom can be used by forming a dative bond with certain metals 152

153 (MC O, Thus (MC O) types of bonds were assumed to be present in metal carbonyls. In the formation of MC O bonds, the electrons are supplied by the molecules of CO and the metal atom is thus said to have zero-valency. The Number of molecules of carbon mono-oxide which can unite with one atom of the metal is controlled by the tendency of the metal atom to acquire the E.A.N. of the next inert gas. For the stable nonnumeric carbonyl. Where Carbonyls E.A.N. = m + 2y = G M = Atomic number of the metal M Y = No. of CO molecules G = At. No. of next inert gas Atomic Number of the metal Number of electron contributed by CO groups E.A.N. Succeeding inert gas Cr(CO) Kr(36) Fe(CO) Kr(36) Ni(CO) Kr(36) Mo(CO) Xe(54) Ru(CO) Xe(54) W(CO) Rn(86) Os(CO) 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 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 - 153

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. X G m 2y X X 1 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. Mn 2 (CO) 10, CO 2 (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 Mn 2 (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 Co 2 (CO) 8. Co 2 (CO) 8 2Co = 2 x 9e - = 18e - 8Co = 2 x 8e - = 16e - Co-Co bond = 1x 2e - = 2e - Co 2 (CO) 8 = 36e Electrons on one Co atom = 18e - 154

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 2p x of oxygen overlap to ii. form a localised bond. 2p y of carbon and 2p y of oxygen overlap to form a π-bond. iii. 2p z of carbon and 2 p z 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 atomcarbon mono-oxide bonds are readily formed in metal carbonyls. It 155

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 156

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) 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 1.50A o which is shorter by 0.32A o in comparison in Ni-C single bond length (=1.82A o ) found in carbonyls. The C-O bond length in this carbonyl has been found to equal to 1.15 A o. Which is larger that the C-O bond length in CO molecule (=1.128A o ) (Fig. 5.3) Fig. 5.3 L Tetrahedral structure of Ni (CO) 4 molecule 157

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 Sp 3 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 3d 10 4S O CO Ni bond is caused by the overlap between the empty sp 3 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 ii. Ni(CO) 2 phen. It confirms that two C=O groups of Ni(CO) 4 are replaced by one molecule of phenanthroline. Similarly the reactions of Ni(CO) 4 with diarsine indicates the two C=O groups are replaced and remaining two are retained. 158

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 1.797A o and 1.842A o respectively. It has dsp 3 hybridisation of Fe atom (Fig. 5.4(c). III. Molecule is Diamagnetic and the distance Fe-C is 1.84A o (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 A o According to old concept when chromium forms Cr(CO) 6 one electron of 4s orbital missing and three 4p orbitals become empty which are hybridised to form six d 2 sp 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. 159

160 When chromium atoms form chromium carbonyl [Cr(CO) 6 ] the metal atom exhibits d 2 sp 3 hybridisation. Out of 6 d 2 sp 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 Mo W

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; e g metal-metal bonding is evident in Mn 2 (CO) 10 as in Figure. 5.6 Structure of Dinuclear Carbonyls (a) (i) Mn 2 (CO) 10 : Its structrue is shown in Fig

162 Fig. 5.6 (ii) Structure of Fe 2 (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 coordinate bond (FeC co-ordinate bond). The presence of Fe-Fe bond is supported by the diamagnetic character of Fe 2 (CO) 9 molecule. Fe- Fe bond distance has been found to be equal to 2.46A o. The terminal C-O bond distances from the structure given in figure. (5.7) 162

163 but equal to 7. The co-ordination number of each Fe atom is not equal to 6 Fig. 5.7: Structure of Fe 2 (CO) 8 Similarly structures of Co 2 (CO) 8 can be represented as in Fig. 5.8 Fig. 5.8: Structure of Fe 2 (CO) 8 (b) Structure of Trinuclear Carbonyls Os 3 (CO) 12 and Ru 3 (CO) 12 possess similar structure (Fig. 5.9a) where as Fe 3 (CO) 12 has a different structure 5.9(b). Os and Ru molecules do not have any bridging CO group (Fig. 5.9 (a)). In 163

164 Fe 3 (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 Structure of Fe 3 (CO) 12 According to old concept each iron atoms gets hybridized trigonal bipyramidally (dsp 3 ). 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 A o and one Fe-Fe bond length is 2.56A o. (Fig. 5.9) Fig. 5.9 Fe 3 (CO) 12 Complex Fig. 5.9 (a) M 3 (CO) 12 Structure 164

165 Fig. 5.9 (b) Structure of Fe 3 (CO) 12 (c) Tetra and Hexanuclear Carbonyls On the similar grounds structures of tetranuclear carbonyls, such as M 4 (CO) 12 [M = Co, Rh, Ir] and hexanuclear carbonyls, M 6 (CO) 16 e g. Rh 6 (CO) 16 can be represented as in Figs and 5.11 respectively. Some heteronuclear carbonyls are also known e.g. Mn 2 Fe(CO) 14 is shown in Fig (a) (b) Fig (a) Structure of Ir 4 (CO) 12 (b) M 4 (CO) 12 ; M = Co or Rh 165

166 Fig Structure of Rh 6 (CO) 16 Fig Structure of Mn 2 Fe(CO) 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 Mn 2 (CO) 10 and bridging carbonyl groups, as in Co 2 (CO) 8. Metal-carbon distances in Fe 2 (CO) 9 and Co 2 (CO) 8 fall into two groups, metal-bridging carbonyl distances being about 0-1 A longer than metal-terminal carbonyl distances. Such a difference is compatible with the concept of two-electron donation by terminal carbonyls, and oneelectron donation (to each of two metal atoms) by bridging carbonyls, through the possible existence of bonds of different strengths makes 166

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 cm -1 ; and Fe 2 (CO) 9 and Co 2 (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 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 cm -1, while bridging carbonyl frequency is obtained in the cm -1 region. While, the strong metal carbon bonding is indicated by very low carbonyl structing frequencies of 1883 and 1788 cm -1 respectively. 167

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 Mn 2 (CO) 10 has...bridging carbonyl group, Fe 2 (CO) 9 has METAL NITROSYLS Nitrosyls are the compounds in which the nitrogen of the nitrosyl group is directly bonded to the atoms or ions, or the compounds containing nitric oxide group are called nitrosyl compound. NO molecule is an odd electron molecule having an unpaired electron, it readily unites with other elements by direct addition to form nitrosyl compounds. Nitric oxide form nitrosyl compound by the following 3 ways: i. A positive ion, NO + is formed due to the loss of an electron which then combine with atom or molecule (:N:::O) + or (:N O:) + 168

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: :N I II III IV V D D :N :N :N :N :O: :O: :O: :O: - :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) 5 NO] Neutral NO and NO - Complexes As has been pointed out metal complexes of neutral NO and its anion, NO - are very rare. Fe(NO) 2 (CO) 2 is supposed to be the important example of metal complex with neutral nitric oxide molecule. This is prepared by the action of nitric oxide on Fe(CO) 5. During the reaction, 169

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)NH 3 ) 5 NO]X 2. 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(NH 3 ) 3 NO] 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) 5 NO] 3- anion is also supposed to be a complex of NO - anion, since it gives NO - stretching frequency at 1150 cm Complexes of NO + Most complexes of nitric oxide and transition metals are best considered to be those of the NO + ion, three electrons being transferred to the metal atom: M-N back π-bonding then takes place in exactly the same way as for carbon monoxide. Because of its positive charge, however, coordinated NO is a better π-acceptor than coordinated CO, and the N-O 170

171 stretching frequency in complexes of NO + is some cm -1 lower than that in salts such as NO + BF - 4. Two NO + derivatives of iron may be mentioned briefly here. The species formed in the brown-ring test for nitrate is [Fe(H 2 O) 5 NO] 2+. The equilibrium [Fe(H 2 O) 6 ] 2+ + NO [Fe(H 2 O) 5 NO] 2+ + H 2 O 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 d 7 complex of Fe 1 and NO +, but the N-O stretching frequency indicates very strong π-bonding and the intense brown colour strongly suggests Fe 1 -NO + charge transfer 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 45 o C temperature on Fe(CO) 5. While M(NO) 4 nitrosyls of Ru and Co are prepared by the same method using Ru 2 (CO) 9 and Co 2 (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). 171

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) 3 CO 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 Mn 2 (NO) 2 (CO) 7 (isoelectronic with Fe 2 (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, (Ph 3 P) 2 Mn 2 (CO) 8, Fe 2 (CO) 9 and Co 2 (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); (Ph 3 P) 2 Mn(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 Nirtosyl Halide Complexes Volatile diamagnetic nitrosyl halides of formula Fe(NO) 3 X are formed by the action of nitric oxide on iron carbonyl halides in the 172

173 presence of finely divided iron as a halogen-acceptor. These readily lose NO to give [Fe(NO) 2 X] 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, CoX 2 + Co + 4NO 2Co(NO) 2 X 4NiI 2 + 2Zn + 8NO 2[Ni(NO)I] 4 + 2ZnI 2 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) 2 Cl 2, which is probably monomeric and tetrahedral. Nitrosyl halides are also formed by some metals which, so far as is known, do not form nitrosyls or nitrosyl carbonyls. Thus molybdenum and tungsten (but not chromium) carbonyls react with nitrosyl chloride: M(CO) 6 + 2NOCl 20 CH Cl 2 2 M(NO) 2 Cl 2 + 6CO Palladium (II) chloride in methanolic solution yields Pd(NO) 2 Cl 2, and nitrosyl halide molecules or anions are formed also by several other transition metals 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 ] H + + NO - 3 [Fe(CN) 5 NO] CO 2 + NH 4 173

174 In the latter process, two successive equilibria are involved: [Fe(CN) 6 ] 4- + NO 2 - [Fe(CN) 5 NO 2 ] 4- + CN - [Fe(CN) 5 NO 2 ] 4- + H 2 O [Fe(CN) 5 NO] OH - 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- +2NO Ba 2+ +3CO 2 + H 2 O 2[Fe(CN) 5 NO] HCN + 3BaCO 3 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) 5 NO 2 ] 4-. Structure of Nitrosyl Co-ordination Compounds: If we compare the electronic structure of NO with CO, it is observed that NO has an additional electron in antibonding π M.O., which may be readily lost to form the nitrosonium ion, NO +. The additional electron present in π molecular orbital of NO can be supplied to metal atom thus increasing its effective number by one unit and neutral No is itself converted into NO + ion. Then, this NO + is coordinated through nitrogen with the metal atom by donating its lone pair to the metal. 1. Cobalt atom may increase its E.A.N. from 27 to 28 by accepting an additional electron from a neutral molecule of NO: Co + NO Co - + NO + cobalt ion may then combine with one NO+ group and 3 CO molecules to form stable compounds Co - 174

175 + NO + + 3CO Co(NO)(CO) 3 In this compound the E.A.N. of Co is, = 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) 2 N lone pair] overlaps suitable vacant hybrid orbital on M ion (sp 3 in tetrahedral or d 2 sp 3 in octahedral) to form ON + M - -bond and the empty π 2 * or π 1 * 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 175

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 cm -1 in nitrosonium salts. Structure of various groups of nitrosyl complexes are shown in Fig to 5.20 Fig Structure of Fe (NO) 4 Fig Structure of Fe(CO) 2 (NO) 2 Fig. 5.15: Structure of [Mn(NO) (CO) 4 ] 176

177 Fig. 5.16: Structure of [Fe(NO) 2 I] 2 Fig. 5.17: Structure of Fe(NO) 3 Cl Fig. 5.18: Structure of [Ni (NO)I] 4 177

178 Fig. 5.19: Red salt of Diethyl Ester of [FeS 2 (NO) 4 ] 2- Fig. 5.20: Anion of Red salt of [Fe 4 (NO) 7 S 3 ] 5.4 DINITROGEN COMPLEXES In 1965, Allen and Senoff obtained salts containing the [Ru(NH 3 ) 5 N 2 ] 2+ cation by the action of hydrazine hydrate on various compounds of tri- and tetrapositive ruthenium, amongst them ruthenium trichloride and ammonium hexachlororuthenate (IV). Thus, these substances (often called nitrogenyl or dinitrogen complexes, to distinguish them from those containing the nitride ion) have been known for only a few years. 178

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 N 2 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, CoCl 2 + Ph 3 P [Ru(NH 3 ) 5 H 2 O] 2+ NaBH 4 (Ph 3 P) 3 CoH 3 EtOH 3 N [Ru(NH 3 3 ) 5 (N 2 )] 2+ N (Ph 3 P) 3 CoH(N 2 ) (b) Another method of preparation of dinitrogen complex is the reaction of coordinated azide: [Ru(NH 3 ) 5 Cl] 2+ N 3 - MeSO NH 3 3 H [Ru(NH 3 ) 5 N 2 ] 2+ (c) Similarly reaction of (Ph 3 P) 2 Ir(CO)Cl with RCON 3 : (Ph 3 P) 2 Ir(CO)Cl + RCON 3 Ir(PPh 3 ) 2 (CO)(Cl)(N 2.NCOR) EtOH CHCl / 3 (Ph 3 P) 2 IrCl(N 2 ) 179

180 (d) Reaction of coordinated NH 3 with HNO 2 : [O s (NH 3 ) 5 (N 2 )] 2+ + HNO 2 [O s (NH 3 ) 4 (N 2 ) 2 ] H 2 O 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 o C without decomposition. Most are rapidly oxidised by air and decompose on heating gently. The orange solid (Ph 3 P) 3 CoH(N + 2 ) shows reversible displacement with hydrogen, ethylene or ammonia. Some of the reactions of (Ph 3 P) 2 IrCl(N 2 ) (yellow solid) are as follows: (Ph 3 P) 2 IrCl(N 2 ) + Ph 3 P (Ph 3 P) 3 IrCl + N 2 (Ph 3 P) 2 IrCl(N 2 ) + HCl (Ph 3 P) 3 IrHCl 2 + N 2 (Ph 3 P) 2 IrCl(N 2 ) + CO (Ph 3 P) 2 Ir(CO)Cl + N 2 Dinitrogen complexes show an asymmetric IR N N stretching frequency in the range Cm -1 (Raman stretching frequency in N 2 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(NH 3 ) 5 N 2 Ru(NH 3 ) 5 ] 4+ and [Sm(N 5 C 5 Me 5 ) 2 I 2 (N 2 )] are shown in Fig and 5.22 respectively. 180

181 Fig. 5.21: Structure of [Ru(NH 3 ) 5 (N 2 )Ru(NH 3 ) 5 ] Fig. 5.22: Structure of [Sm(N 5 C 5 Me 5 ) 2 I 2 (N 2 )] Fixation of Nirtrogen Dinatrogen complexes while show possibility of developing new methods for nitrogen fixation, they also help in the understanding of the probable mechanism of biological fixation of nitrogen. An important enzyme-system is related with the atmospheric fixation of nitrogen; which involves an important step in nitrogen-cycle and is responsible for supply of nitrogen to the plants growth (e.g. Bluegreen algae, symbiotic bacteria legume) 181

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 and contains Fe 4 S 4 group; while the large protein is a tetramer of molecular weight It has 2 molybdenum atoms, nearly 30 iron atoms and nearly 30 mobile sulphide ions. Fe-S group probably functions as redox centre, and the active site for dinitrogen binding is probably molybdenum atom. (Fig. 5.23) Fig. 5.23: Fixation of Nitrogen 5.5 DIOXYGEN COMPLEXES Amongst all the donor atoms oxygen is most important. The donor ability of oxygen is related with its partial charge; higher is the negative charge, higher will be the donar ability. Large number of coordination compounds are available in which oxygen uses one of its two lone pairs of electrons. The most important example of dioxygen complexation is transportation of oxygen in aerobic-organisms through heme and hemocynin mechanism. Although, hemoglobin and hemocynin are known since long time for their specific ability of absorption and release of 182

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 (Fe 2+ ), 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 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: Hemoglobin Globin Heme Protoporphyrin Fe(II) 183

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 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: I st step: Bonding with dioxyen: II nd step: Bonded dioxygen links with other heme ( -peroxo complex is formed) : III rd step: Decomposition of per oxo complex into ferryl complex. IV th step: Reaction of ferryl Complex with heme to give Hematin: In living being steps I and IV do not take place, otherwise total heme would have precipitated as hematin. Apart from other reactions, 184

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 + 4O 2 Hb (O 2 ) 4 Hemoglobin Oxyhemoglobin While in tissues, the reaction that takes place is: Hb(O 2 ) 4 + 4Mb 4Mb(O 2 ) + Hb Myoglobin 5.6 TERTIARY PHOSPHINE AS LIGAND Large number of triphenylphosphine and similar substituted metal carbonyls are known, e.g. Ni(CO).(Ph 3 P) 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 Ph 3 P in 1:1 ratio of Co and Ph 3 P. One compound is [Co 2 (CO) 6 (PPh 3 ) 2 ] and the other is the salt [Co(CO) 3 (PPh 3 ) 2 ][Co(CO) 4 ] in which the cation has the expected trigonal pyramidal structure. A platinum complex, Pt (CO) 2 (PPh 3 ) 2 can be obtained by the action of CO on Pt(PPh 3 ) 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, (Ph 3 P)Co(CO) 3 I can be made by the remarkable reaction. [(Ph 3 P)Co(CO) 3 ] - 2CF 3 I (Ph 3 P)Co(CO) 3 I + I - + C 2 F 6 Triphenylphosphine carbonyl halides of rhodium and iridium may be prepared by interaction of the metal halide (or a complex halide) and 185

186 triphenylphosphine in a variety of organic solvents, the solvent serving as the source of the carbonyl group: (NH 4 ) 2 IrCl 6 IrCl 3.3H 2 O P Ph 3 (Ph 3 P) 2 Ir(CO)CI P Ph 3 (Ph 3 P) 2 Ir(CO)CI The product of these reactions - (Vaska's compound) is a highly reactive complex. Vaska's compound is a carbonyl halide; and many triphenylphosphine complexes containing rhodium and iridium show similar reactivity and catalytic activity. The iridium compounds is remarkable for its reversible uptake of H 2, O 2 and SO 2 to give crystalline 1:1 adducts which can be decomposed by lowering the pressure; for example, (Ph 3 P) 2 Ir(CO)Cl O O 2 2 Ir(PPh 3 ) 2 (CO)Cl In the oxygen adduct, oxygen atoms occupy cis octahedral positions; the O-O distance of 1.30 A o suggests that Oxygen is present as O 2 - rather than O Some of the many other reactions of Vaska's compounds are shown in Fig 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, O 2, N 2 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 V 2 (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 cm -1 (or cm -1 ). While the bridging CO - group is indicated by the stretching frequency at 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- 188

189 :N (a) (b) (c) (d) D D :N :N :N :O: :O: :O: :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, (Ph 3 P) 2 Mn(CO) 8, Fe 2 (CO) 9 and Co 2 (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 Mn 2 (NO) 2 (CO) 7 are also formed. Many dinitrogen complexes have been reported e.g. [Ru(NH 3 ) 5 N 2 ], [(NH3) 5 RuN 2 Ru(NH 3 ) 5 ], [(Ph 3 P) 2 IrCl(N 2 )], [O s (NH 3 ) 4 (N 2 ) 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 ) protein contains Fe 4 S 4 groups; while the large protein (mol wt ) is a 2 2 tetramer, which has 2 Mo atoms, nearly 30 Fe atoms and nearly 30 mobile sulphide ions. Fe-S group functions as 189

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 + 4O 2 Hb(O 2 ) 4 Hemoglobin Oxyhemoglobin Hb(O 2 ) 4 + 4Mb 4Mb(O 2 ) + Hb Myglobin Oxyhemoglobin Large numbers of triphenyl phosphine and similar substituted metal carbonyls are known e.g. Ni(CO)(Ph 3 P) 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' (Ph 3 P) 2 Ir(CO)Cl, which is used for the preparation of large number of triphenyl phosphine containing complexes. 190

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 cm -1 and at cm -1 at 1883 and 1788 cm -1 respectively (d) Has no bridging group, Fe 2 (CO) 9 has three bridging groups. 2(A) (i) With NO+ ion (ii) :N formula M(NO) 4 (M = Fe, CO and Ru). (a) (b) (c) (d) D D :N :N :N :O: :O: :O: :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 contains Fe 4 S 4 group Large protein mol. wt contains 2 Mo atoms 30 Fe atoms and 30 mobile sulphide ions. C.(i) To give Oxohemoglobin Hb + 4O 2 Hb(O2) 4 (ii) Myoglobin for stroage and (iii) by cytochromes D. Vaska compound has general formula: (Ph 3 P) 2 Ir(CO)Cl 191

192 Unit - 6 REACTION MECHANISM OF TRANSITION METAL COMPLEXES-I Structure 6.0 Introduction. 6.1 Objectives. 6.2 Energy Profile of a Reaction Reactivity of metal Complex - Inert and Labile Complexes Valence Bond and Crystal Field applications. 6.3 Kinetics of Octahedral Substitution Nucleophilic Substitution Hydrolysis Reactions Factors affecting Acid Hydrolysis Base- Hydrolysis-Conjugate Base Mechanism Anation Reaction Reactions without Metal-Ligand Bond-Cleavage 6.4 Let Us Sum Up 6.5 Check Your Progress: The Key 192

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 193

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 SN 1 and SN 2 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 194

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 following relation: G = H - T S, and the disorder, S, according to the 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 Reactivity of Metal Complex - Inert And Labile Complexes Almost all the reactions of transition metal complexes may be divided into two categories; (a) Substitution reactions and (b) Redoxreactions. In coordination chemistry rate of a reaction is equally important, as the reaction equilibrium. The ability of a complex to engage itself in reactions involving the replacement of one or more ligands in its coordination sphere by other 195

196 ligands is called the lability of the complex. The complexes that undergo rapid substitution (half time period or T 1/2 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 14 CN - 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(H 2 O) 6 ] 3+ (bond energy = 690 KJ mol -1 ) is labile while [Cr(H 2 O) 6 ] 3+ is inert. [Ni(CN) 4 ] 2- is thermodynamically stable but kinetically labile but [Co(NH 3 ) 6 ] 3+ is kinetically inert but thermodynamically unstable 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 sp 3 d 2 hybridisation. ii. Inner-Orbital Complexes which involve d 2 sp 3 hybridisation. The two d-orbitals involved in sp 3 d 2 and d 2 sp 3 hybridisation are dx 2 -y 2 and dz 2 eg set orbitals. 1. Outer-Orbital Octahedral Complexes Outer-orbital octahedral complexes (sp 3 d 2 hybridisation) are generally labile for example the octahedral complexes of Mn 2+ (3d 5 )Fe 2+, Fe 3+ (3d 5 ) Co 2+ (3d 7 ) Ni 2+ (3d 8 ) Cu 2+ (3d 9 ) and Cr 2+ (3d 4 ) exchange ligands rapidly and hence are labile. This is 196

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 d 2 sp 3 hybrid orbitals are filled with the six electron pairs denoted by the 6 ligands. d n 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 t 2g 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 t 2g set (i.e. dxy, dyz and dxz) contains at least one electron, and have no vacant orbtial to link an extra ligand. 197

198 Type of the complex inner orbital labile octahedral complexes inner orbital inert octahedral complexes Table 6.1 Distribution of d n -electrons in various t 2 g orbitals for labile and inert inner-orbital octahedral complexes (according to VBT) d n configu ration Distribution of d n electron (shown by arrows) in t 2 g orbitals. Electrons shown by crosses in e g orbitals have been donated by six ligands to enter d 2 sp 3 hybrids and are in opposite spins. Example of central metal ions d s p t 2 g e g xy yz zx x 2 -y 2 z 2 px py pz d 0 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) d 1 xx xx xx xx xx xx Ti(+3), V(+4), Mo (+5), W Re(+6) d 2 xx xx xx xx xx xx Ti(+2), V(+3), Nb (+3), Ta(+3), W(+4), Re(+5), Ru(+6) d 3 xx xx xx xx xx xx V(+2), Cr(+3), Mo(+3), W(+3), Mn(+4), Re(+4) d 4 xx xx xx xx xx xx [Cr(CN) 6 ] 4-, Mn(CN) 6 ] 1 Re(+3), Os(+3), Ir(+4) d 5 xx xx xx xx xx xx [Mn(CN) 6 ] 4-, Re(+2), Fe(CN) 6 ] 3- Ru(+3), Os(+3), Ir(+4) d 6 xx xx xx xx xx xx [Fe(CN) 6 ] 4-, Ru(+2), Os(+2), Co(+3) 4- (except Co Fe 3 Rh (+3), Ir (+3) 198

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 S N 2 associated mechanism SN2 association mechanism Octahedral (oct.)pentagonal bipyramidal (pent.bipy.) (C.N. = 6) (C.N. = 7) d n ion Strong ligand field (spin-paired or low-spin complexes) Weak ligand field (spin-paired or low-spin complexes) Oct. pent.bipy. E Oct. pent.bipy. E (C.N.= 6) (C.N.= 7) (C.N.= 6) (C.N.= 7) d 0 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq d d d d d d d d d d 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 S N 1 associated mechanism 199

200 SN1 association mechanism Octahedral (oct.) Syware Pyramidal (Squ. pyi) (C.N. = 6) (C.N. = 5) d n ion Strong ligand field (spin-paired or low-spin complexes) Weak ligand field (spin-paired or low-spin complexes) Oct. pent.bipy. E Oct. pent.bipy. E (C.N.= 6) (C.N.=5) (C.N.= 6) (C.N.= 5) d 0 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq 0 Dq d d d d d d d d d d 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 (S N 1, and S N 2). Negative values of E denotes a loss of CFSE when octahedral complex is changed into an activated complex which may be square pyramidal or pentagonal bipyramidal. If the CFSE of the activated complex is greater than that of octahedral complex. E has been given zero value which shows that these complexes do not loose CFSE when they are changed into activated complexes. 200

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 d 0, 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 d 3, d 4, d 5, and d 6 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 SN 1 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 d 3, d 4, d 5 and d 6 ions is: Order of inertness : d 6 > d 6 > d 4 > d 5 Loss of CFSE for SN 1 mechanism : -4.00>-2.00>-1.43>-0.86 Loss of CFSE for SN 2 mechanism : The order of reactivity will be reverse of the above i.e. the order of reactivity will be d 6 > d 3 > d 4 > d 5 it is supported by the following facts: i. High spin octahedral complexes of d 3 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 SN 1 or SN

202 ii. High spin octahedral complexes of d 5 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 d 8 ion are inert. According to VBT d 8 ion [3dxy 2, 3yz 2, 3dxz 2, 3d(x 2 -y 2 ), 3dz 2 ] will form outer orbital complexes and will be labile. Therefore in case d 8 ion VBT & CFT gives different predictions. iv. Both high spin and low spin octahedral complexes of d 10 ion are labile. Factors Affecting the Liability of Complex 1. Charge of the metal ion: For the isoelectronics complexes there is a decrease in lability with the increase of the charge of the central metal ion. i. The order of lability of the complex is as follows: Lability order : [AlF 6 ] 3- > [SiF 6 ] 2 > [PF 6 ] - > [SF 6 ] 0 Cationic charge : +3 < +4 < +5 < +6 ii. The rate of water exchange represented by: [M(H 2 O) 6 ] n + 6H 2 O * [M(H 2 O * ) 6 ] n+ + 6H 2 O decreases with the increase of cationic charge in the series Rate of water exchange: [Na(H 2 O) 6 ] + > [Mg(H 2 O) n ] 2+ > [Al(H 2 O) 6 ] 3+ Cationic Charge +1 < +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(H 2 O) 6 ] 2+ <[Ca(H 2 O) 6 ] 2+ < [Sr(H 2 O) 6 ] 2+ Cationic Size (A) 0.65 < 0.99 <

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(H 2 O) 6 ] 2+ (a d 8 system) has the largest value of half life i.e. it reacts slowest. The hydrated M 2+ ions [M(H 2 O) x ] 2+ of the first row transition elements are all high spin complexes. ii. [Cu(H 2 O) 6 ] 2+ reacts most rapidly because the 2 water molecules above and below the square plane of the tetragonal distorted octahedral shape of [Cu(H 2 O) 6 ] 2+ are exchanged. The remaining four H 2 O molecules lying in the square plane react slowly. Fig. 6.2: Half-lives (in sec) at 25 o C 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 ] CN - [Ni( 14 CH) 4 ] CN 203

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 5 th group into the coordination sphere to form on activated complex. Check Your Progress-1 Notes : (i) Write your answers in the space given below. (ii) Compare your answers with those given at the end of the unit. A.(i)For a reaction to go in the forward direction G should be... (ii)according to the thermodynamic relation G =... That is for conversion of the reactants into the products, the bond energies and the state of disorder should be...i.e. the value of the H should be...and that of T S should be... B.(i) According to VBT, generally labile complexes are... complexes, while the inert complexes are...complexes. (ii) Inner orbital complexes may be labile, if they have at least... in...set is vacant, e.g. in... (iii) According to CFT inert complexes have...values of

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) SN 1 and SN 2 1. In SN 1 process the rate-determining slow step is a metal-ligand bond breaking step, since the co-ordination No. of the complex MX 5 Y (=6) is decreased to 5 which is the co-ordination number of the intermediate MX 5. For a ligand replacement reaction of the general type [L n MX] + Y = [L n MY] + X (For simplicity all charges are omitted), the mechanism analogous to unimolecular nucleophilic substitution (sn 1 ) at a carbon atom would be: [L n MX] slow [L n M] + X [L n M] + y fast [L n MY] + X The rate of SN 1 mechanism is first order with respect to MX 5 Y, i.e. the rate-determining step in this mechanism is unimolecular. On the other hand the rate determining step for SN 2 mechanism is bimolecular, i.e. its rate of reaction is second order 205

206 first order with respect to MX 5 Y and first order with respect to Z. Thus for SN 1 mechanism rate = K [MX 5 Y], and for SN 2 mechanism rate = K [MX 5 Y][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: This leads to [L n MX] + Y k 1 [L n MX] Y [L n MX] y slow [L n MY] + X where d [Ln MY] = dt k 1 k2[ LnMX ][ Y] k 1 k = k[l n MX][Y] 2 k = k1k2 k k 1 2 Detailed investigation of such a reaction can lead to a value for k 1 /k -1, the equilibrium constant for ion-pair formation Nucleophilic Substitution As has been motioned, nucleophile substitution reactions in octahedral complexes follow either of the two mechanisms, the dissociation mechanism or the SN - 1 mechanism and the association mechanism or the SN - 2 mechanisms. The rate determining step in association or dissociation, may be worked out by analysing the rate-laws of the reactions taking place and the specific conditions under which the reactions take place. The difference in these two mechanisms depends on, 206

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, MX 5 Z. There two steps are diagrammatically shown in Fig. 6.3 Fig. 6.3: S N 1 or dissociation mechanism for the substitution reaction MX 5 Y + Z MX 5 Z + Y For the S N 1 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 S N 1 substitution (k 1 ) should be inversely proportional to the strength of the Co-L bond, and depend on the charge, steric factors, and chelating effects of the leaving group L. 207

208 (3) Increase in the electron density on Co atom by the electron donors in S Nn should assist the M-L bond breaking. (4) k 1 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 (NH 3 ) or are themselves acceptors (NO - 2, 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 d 3, d 6 and d 8 and high spin d 3 and d 8 ions. Hence, these complexes do not favour the S N 1 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. 208

209 System High spin Low spin CN = 5 CN = 7 CN = 5 CN = 7 d o, d d d d d d d d d d 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 MX 5 Y 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. MX 5 Y (C.N.=6) ( Z ) Slow MX 5 YZ Unstable seven-coordinatee Intermediate (C.N.=7) This reaction is rate-determining and bimolecular because two reactants viz MX 5 Y and Z are involved in this step. Thus the rate of this rate-determining reaction is of second order: first order 209

210 with respect to the complex, MX 5 Y and first order with respect to the entering ligand, Z, i.e., Rate of reaction = K[MX 5 Y][Z] In the second step either at the same as Z adds to MX 5 Y or shortly thereafter, Y leaves MX 5 YZ rapidly to give MX 5 Z. This is a fast step. MX 5 YZ Fast MX 5 Z 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-MX 5 -Z] in the pre-equilibrium step: Thus the following equilibrium will be established: LMX 5 + Z MX 5. Z; K = [ LMX [ LMX 5 5. Z] ][ Z] The value of the equilibrium constant, K, for the association complex, may be obtained using Fuoss-Eagan equation, where, K = 3 4 a 3 N A e -v/rt a = Nearest reach-distance v = Coulomb potential energy at a-distance N A = Avogadro number = (Z 1 Z 2 e 2 /4 ea) As in the octahedral complexes, the six ligands are already present along the three C 4 axes along which the e g orbitals are concentrated, the t 2g orbitals that lie along the C 2 axes most probably have to be approached by the seventh ligand to form the associated complexes in the S N 2 process. 210

211 Hence, if the t 2g orbitals are filled (Co 2+ in low spin octahedral complexes), the higher activation energy required to empty one of the t 2g orbitals will make the complex inert. Table 6.4 also shows that due to the loss of the CFSE energies, the d 3 and low spine d 6 ions require highest activation energies, followed by d 7, d 8 (Ni 2+ complexes are labile due to the expulsion of ligand by the e g orbitals) and high spin d 3 and d 8 ions. Thus, S N -1 and S N -2 reactions differ in the following points: (i) In S N 1 process the rate-determing slow step is a metal-ligand bond breaking step, since the coordination number of the complex, MX 5 Y (=6) is decreased to 5 which is the coordination number of the intermediate, MX 5. On the other hand in S N 2 process the rate-determining step involves a metal-ligand bond making step, since C.N.=6 is increased to 7. (ii) The rate of S N 1 mechanism is first order with respect to MX 5 Y, i.e., the rate-determining step in this mechanism is unimolecular. On the other hand the rate-determining step for S N 2 mechanism is bimolecular, i.e. its rate of reaction is second order: first order with respect to MX 5 Y and first order with respect to Z. Thus: for S N 1 mechanism rate = K[MX 5 Y] and for S N 2 mechanism rate = K[MX 5 Y][Z] Hydrolysis Reactions The substitution reactions in which a ligand is replaced by a H 2 O 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 H 2 O molecules are called acid hydrolysis or equation reactions, 211

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 (ph 10). Examples are:- [Co III (NH 3 ) 5 Cl] 2+ + H 2 O[Co III (NH 3 ) 5 (H 2 O)] 3+ + Cl - [Co III (en) 2 ACl] + + H 2 O [Co III (en) 2 A(H 2 O)] 2+ + Cl - [A = OH -, Cl -, NC -, NO 2 - ] [Co(NH 3 ) 5 Cl] 2+ +OH - [Co(NH 3 ) 5 (OH)] 2+ +Cl - (Base hydrolysis reaction) (a) Acid Hydrolysis or Aquation : When NH 3 or ammines like ethylene diamine or its derivatives co-ordinated Co(III) are replaced very slowly by H 2 O 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(NH 3 ) 5 X] 2+ + H 2 O[Co(NH 3 ) 5 (H 2 O)] 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. Acid hydrolysis reaction The rate law K = K 1 [Co(NH 3 ) 5 X] 2+ [55.5] does not indicate whether these reactions proceed by an SN 2 displacement of X by H 2 O or by an SN 1 dissociation followed by the addition of H 2 O. 212

213 The rate law for acid hydrolysis at low ph thus becomes d - [Co(NH3 ) 5 X] = k A [Co(NH 3 ) 5 X] dt (If X is the anion of a weak acid, a term k H +[Co(NH 3 ) 5 X][H + ] is added.) As we have shown previously, such a rate law is compatible with either a slow dissociation of the complex into [Co(NH 3 ) 5 ] 3+ and X or replacement of X by H 2 O 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) 2 Cl 2 ] +, where AA is a substituted ethylendiamine, were examined. For replacement of a single chloride ion at ph 1 the order found for values of k A was CH 2 NH 3 CH 3 CHNH 2 CH 3 CHNH 3 (CH 3 ) 2 CNH 2 < < < CH 2 NH 3 CH 2 NH 2 CH 3 CHNH 2 (CH 3 ) 2 CNH 2 Such an acceleration of substitution by bulky ligands suggests that the dissociative mechanism is operative; although introduction of methyl groups must have some inductive effect, the variation in base strengths among the diamines is very much less than the variation in rate constants for the hydrolysis of their cobalt (III) complexes, and it seems reasonable to attribute the kinetic effect mainly to steric factors. Now since steric factors favour S N 1 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(NH 3 ) 5 X] and the formation constant of the [Co(NH 3 ) 5 X] complex from [Co(NH 3 ) 5 (H 2 O)] 3+ and X, and 213

214 (b) the decrease in the rate of the exchange reaction [Co(NH 3 ) 5 (H 2 O)] 3+ H 18 2 O = [Co(NH 3 ) 5 (H 18 2 O)] 3+ + H 2 O at high pressures 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(NH 3 ) 4 (H 2 O)Cl 2 ] 2+ occurs in two steps: [Co(NH 3 ) 4 (H 2 O)Cl] 2+ + slow Cl [Co(NH 3) 4 (H 2 O)] 3+ + Cl - 6-co-ordinate complex 5-co-ordinate Intermediate [Co(NH 3 ) 4 (H 2 O)] 3+ + fast H 2O [Co(NH 3) 4 (H 2 O) 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 NH 3 molecules in [Co(NH 3 ) 5 Cl] 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 -CH 2 -CH 2 or -(CH 2 ) 2 -chelated links increases the rate values decreases. 214

215 The replacement of NH 3 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) 2 Cl 2 ] + are replaced by the alkyl groups like CH 3,C 2 H 5 etc. the ligand becomes more bulky, if the strained complex having bulky ligand reacts by SN 1, dissociative mechanism and co-ordination number 6 is reduced into 5 co-ordinated intermediate, on the other hand if the strained complex reacts by SN 2 displacement process, the crowding on the complex is increased as it is converted into a transition state of coordination number seven. The rate of hydrolysis of trans [Co(AA 2 Cl 2 )] + at 25 o C and ph=1 corresponding to the replacement of only one Cl- ion by H 2 O molecule are given. Here AA is the diamine. (iv) Effect of Leaving Group The rate of reaction of [Co(NH 3 ) 5 X] 2+ corresponding to the replacement of X with H 2 O 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. HCO - 3 >NO - 3 >I - >Br - >Cl - >SO -- 4 >F - >CH 3 COO - >SCN - <NO Base- Hydrolysis-Conjugate Base Mechanism The base hydrolysis reaction represented by following equation: [Co(NH 3 ) 5 Cl] 2+ + OH - [Co(NH 3 ) 5 (OH)] 2+ + Cl - It involves following two mechanisms. 215

216 1. SN -2, Displacement Mechanism [Co(NH 3 ) 5 Cl] 2+ The reaction proceeds as: ) slow ( OH [Co(NH 3 ) 5 (OH)Cl] + fast [Co(NH 3 ) 5 (OH)] 2+ + Cl - (C.N.=6) (C.N.=7) (C.N.=6) Rate of Reaction = K[Complex][base] = K[Co(NH 3 ) 5 Cl][OH - ] d The rate law is - [Co(NH3 ) 5 Cl] = K B [Co(NH 3 ) 5 Cl][OH - ] dt 2. SN -1 Displacement Mechanism: The complex which acts as a Bronsted acid is converted into its conjugate base (CB), [Co(NH 3 ) 4 (NH 2 )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 k n is nearly 10 4 times higher than k A. Why should hydroxyl ions posses the exceptionally high nucleophilic activity as compared to the similar anions? (3) If a proton cannot be abstracted from N 5 ligands (e.g., [Co(py) 4 - Cl 2 ] + or [Co(CN) 5 Cl] 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 N 5 group giving CB of the ligand. This undergoes the dissociative mechanism as shown below: [(NH 3 ) 5 CoCl] 2+ +OH - fast [(NH 3 ) 4 Co(NH 2 )Cl] + +H 2 O (6.1) [(NH 3 ) 4 Co(NH 2 )Cl] + slow [(NH 3 ) 4 Co(NH 2 )] 2+ + Cl - (6.2) [(NH 3 ) 4 Co(NH 2 )] 2+ +H 2 O fast [(NH 3 ) 5 Co(OH)] 2+ (6.3) 216

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 S N 1CB process. The rate law will be- d = [Co(NH3 ) 5 OH] = dt where, K = K1K 2[ Co( NH3) 5Cl ][ OH] 2 K1[ H2O] K2[ H2O] = K[Co(NH 3 ) 5 Cl][OH] K 1 K1K2 2 [ H2O] K2[ H2O] Though it seems very unlikely that reduction in one positive charge form [Co(NH 3 ) 6 ] 3+ to [Co(NH 3 ) 5 (NH 2 )] 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 NH 2 group. The S N 1 CB mechanism does not explain the following observations. (i) The conjugate base readily dissociates and releases the ligand L; and (ii) the concentration of the conjugate base is very low due to the basic nature of the ligands, and should be present only as a very small fraction of the concentration of the complex present. 217

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 H 2 O and conjugate base, [Co(NH 3 ) 4 (NH 2 )Cl] + of [Co(NH 3 ) 5 Cl] 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) 2 Br] and trans [Co-(Py) 4 Cl 2 ] + 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 SN 1 CB mechanism is not possible. Such complexes undergo rapid base hydrolysis supports the SN 1 CB mechanism and the acid-base properties of the complexes are more important to the rate of reaction, than the nucleophilic properties of OH. Thus both the mechanisms give the same rate laws and the same hydroxo products in aqueous solution, because water is a good cocoordinating agent. The rate of formation of [Co(en) 2 (NO 2 )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 N - 3, NO - 2, SCN - ion are used as nucleophiles. In SN 1 CB mechanism the reactions of [Co(NH 3 ) 5 Cl] 2+ and OH - in aqueous solution at 25 o C in presence of H 2 O, when H 2 O 2 is added to the reaction mixture of [Co(NH 3 ) 5 Cl] 2+ and OH -, the reaction between OH - and H 2 O 2 occurs as: OH H 2 O 2 H 2 O + HO 2 218

219 Which increase the rate of base hydrolysis reaction and form peroxo products. On the other hand if the reaction occurs by an SN 1 CB mechanism the addition of H 2 O 2 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 1 CB reaction is directly proportional to the concentration of OH Anation Reaction The reaction in which an aquo ligand (i.e. H 2 O molecule) from an aquo complex is replaced from the co-ordination shell by some axion. [Co(NH 3 ) 5 (H 2 O)] 3+ + X - [Co(NH 3 ) 5 X] 2+ + H 2 O 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(NH 3 ) 5 (H 2 O)] 3+ slow [Co(NH 3 ) 5 ] 3+ fast [Co(NH 3 ) 5 X] 2+ + H 2 O Let us consider replacement of water in a species containing five non-labile ligands such as [Co(NH 3 ) 5 (H 2 O)] 3+, and let us reverse the experimental procedure and attempt to infer kinetic behaviour from a postulated mechanism. This is k [L 5 M(H 2 O] 1 [L 5 M] + H 2 O k [L 5 M] + Y 1 [L 5 MY] Since Y competes with solvent water for the active intermediate [L 5 M], the rate of formation of [L 5 MY] can be dependent on the concentration of Y. On the other hand, there should be some high 219

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 [L 5 MY] at this concentration should be equal to the rate of formation of [L 5 M] and also equal to the rate of exchange of water between [L 5 M(H 2 O)] and the solvent. Thus the rate of formation of [L 5 M] is given by d [L5 M] = k 1 [L 5 M(H 2 O)] - k -1 [L 5 M][H 2 O] - k 2 [L 5 M][Y] dt According to the steady-state approximation, the concentration of the very reactive [L 5 M] remains small and constant during the reactions d (i.e. [L5 M] = 0 at the steady state). Thus, dt and [L 5 M] = K1[[L5M( H 2O] K [ H O] K [ Y] 1 d [L5 MY] = dt if k -1 [H 2 O] > k 2 [Y] d [L5 MY] = dt 2 2 K1K2[L5M( H 2O][ Y] K [ H O] K [ Y] 1 K K 1 2 K 1 2 [L5M( H 2O][ Y] 2 and a second-order reaction will be observed. On the other hand, if k 2 [Y] > k -1 [H 2 O] d [L5 MY] = k 1 M( H ] dt [L 2 5 O giving first-order kinetics with the overall first-order constant equal to that for the dissociations of the aquo complex 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(NH 3 ) 5 H 2 O] 3+, from carbon a to complex, 220

221 [Co(NH 3 ) 5 CO 3 ] and nitrito complex [Co(NH 3 ) 5 ONO] 2+ from hydroxo or aquo-complex, [Co(NH 3 ) 5 H 2 O] 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 H 2 O (Fig. 6.6). Similarly the reaction of pentamineaquacobalt (III) with NO - 2 ion is explained by the sequence in Fig Fig. 6.6 Mechanism of substitution of carbonate group by water through electrophilic attack by H 2 O +. Fig. 6.7 Probable mechanism of substitution of [(NH 3 ) 5 Co(OH)] 2+ by nitrite NO 2 - through an electrophilic attack. 221

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) S N -1 and S N -2 mechanism of substitution reactions differ in - (i) In S N -1 rate determining step is...process, while that in S N -2 is...process. (ii) The rate determining step in S N -1 is...molecular, while that in S N -2 is... molecular. (iii) In S N -1 rate =... while that in S N -2 rate =... (b)(i) For acid hydrolysis at low ph, the rate Law is -...=... (ii) The rate law of base hydrolysis reactions of an octahedral ammine complex, by SN -1 CB process is - d = [Co(NH3 ) 5 OH] =... dt =... (iii) Formation of aquo-complex from a carbonato complex is an example of substitution... bond, and involves... attack on LET US SUM UP In order to convert reactions in to products it is necessary that the groups or the atoms linked in what ever manner in the reactants molecules should separate and then reunite in the form of the products. 222

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 - T S, hence the possibility of conversion of reactants in to products is only when the state of disorder, and the bond energies in the products, are relatively high; H should be negative and S should be positive. Complexes are generally classified as labile and inert with reference to their reactivity. The ability of a complex to engage itself in the reactions involving the replacement of one or more ligands in the coordination sphere by other ligand is called lability of the complex. The complexes that undergo rapid substitution are termed labile; where as these with law rates of substitution are called inert. According to VBT, the inner orbital complexes (using d 2 sp 3 hybridisation for octahedral complication) are inert while the outer orbital complexes (using sp 3 d 2 hybridisation) are labile. The inner orbital complexes may be labile only when they have at least one d-orbital in t 2g set is vacant; e.g. [V(NH 3 ) 6 ] 3+ ion, According to CFT, complexes with high values of CFSE are inert, while those with small values of CFSE are labile. Substitutions of ligands generally follow one of the two, S N -1 or S N -2 mechanisms. 223

224 In S N -1 or dissociation mechanism, the rate determining slow step is a metal-ligand bond breaking step, since the coordination number of the complex, MX 5 Y (=6) is decreased to 5 in the intermediate, MX 5, complexes: [MX 5 Y] Y [MX 5 ] Z [MX 5 Z] (C.N.=6) (C.N.=5) (C.N.=6) Thus, the rate of SN -1 mechanism is first order with respect to MX 5 Y, i.e. the rate determining step is unimolecular. In S N -2 process the rate determining step involves a metal-ligand bond making step, with the increase of coordination number from 6 to 7: [LMX 5 ] Z [L-MX 5 Z] Y [MX 5 Z] (C.N.=6) (C.N.=7) (C.N.=6) Thus, the rate determine step in S N -2 process in bimolecular i.e. its rate of reactions is second order; first order with respect to [MX 5 L] and first order with respect to Z. For S N -1 mechanism, rate = k [MX 5 L] For S N -2 mechanism, rate = k [MX 5 L][Z] The substitution reaction in which a ligand is replaced by a H 2 O 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 H 2 O molecule while the reaction in which a hydroxo complex is formed by the replacement of a ligand by - OH group is called base-hydrolysis. Acid hydrolysis occur, in neutral and acid solutions (ph<3), while base hydrolysis occurs in basic solutions (ph>10). 224

225 The rate of acid hydrolysis reaction is of first order - [Co (NH 3 ) 5 X] 2+ + H 2 O [Co(NH 3 ) 5 (H 2 O)] 3+ + X - As in oqueous solutions the concentration of water is always constant, the rate law, K = K 1 [Co (NH 3 ) 5 X] 2+ [55.5] does not indicate whether these reactions proceed by and S N -2 displacement of X by H 2 O or by S N -1 dissociation followed by the addition of H 2 O. For base hydrolysis; [Co (NH 3 ) 5 Cl] OH [Co(NH 3 ) 5 (OH)] 2+ + Cl - SN-2 mechanism gives rate of the reaction = K[Co(NH 3 ) 5 Cl] 2+ [OH - d ] and the rate law - [Co(NH3 ) 5 Cl] = K B [Co(NH 3 ) 5 Cl][OH - ]. dt Gerick proposed S N -1CB mechanism for base hydrolysis reaction. In this, the - OH ions abstract a proton from a ligand in N 5 group, giving the conjugate base of the ligand. This under goes the dissociative mechanism: [(NH 3 ) 5 CoCl] 2+ + OH - [(NH 3 ) 4 Co(NH 2 )Cl] + [(NH 3 ) 4 Co(NH 2 )] 2+ + H 2 O Fast [(NH 3 ) 4 Co(NH 2 Cl)] + + H 2 O Slow [(NH 3 ) 4 Co(NH 2 )] 2+ + Cl - Fast [(NH 3 ) 5 Co(OH)] 2+ Thus, the rate determining step is the dissociation of the amido group. The rate law will be- d [Co(NH3 ) 5 OH] = dt - K1K 2[Co(NH3 ) 5Cl][OH ] 2 K1[ H2O] K2[ H2O] = K [Co(NH 3 ) 5 Cl][OH - ] The reactions involving removal of coordinated water molecule are known as 'anation' reactions: [[Co(NH 3 ) 5 (H 2 O)] 3+ + X - [Co(NH 3 ) 5 X] 2+ + H 2 O 225

226 This reaction is reverse of acid hydrolysis reaction. The same second order kinetics would be observed for a unimolecular process: [Co(NH 3 ) 5 (H 2 O)] 3+ Slow H 2O [Co(NH 3) 5 ] 3+ Fast X [Co(NH ) X] Many a times replacement of ligand takes place without breaking a metal-ligand bond, e.g. formation an aquo-complex from a carbonato complex. These involve the electrophilic attack by the proton on the O-atom bonded to the metal. 6.5 CHECK YOUR PROGRESS: THE KEY 1(a)(i) G should be negative. (ii) G = H - T S That is...should be very high i.e. H should be negative and T S should be positive. (b)(i) Labile complexes are outer orbital complexes...inert complexes are inner orbital complexes. (ii) One orbital in t 2g set: e.g. [V(NH 3 ) 6 ] 3+ (iii) Have high values of CFSE. 2.(a)(i) Is metal-ligand bond breaking process that in S N -2 is metalligand bond making process. (ii) S N -1 is Unimolecular. S N -2 is bimolecular. (iii) S N -1, rate = K[MX 5 Y] S N -2, rate = K[MX 5 Y][Z] 226

227 (b)(i) Rate law is - d - [Co(NH3 ) 5 X] =K A [Co(NH 3 ) 5 X] dt (ii) = - K1K 2[Co(NH3 ) 5Cl][OH ] 2 K [ H O] K [ H O] = K [Co(NH ) Cl][OH ] 3 5 (iii) Without breaking M-L bond. Involves electrophilic attack of proton on oxygen bonded with metal. 227

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 228

229 UNIT 7 REACTION MECHANISM OF TRANSITION METAL COMPLEXES-II Structure 7.0 Introduction 7.1 Objectives 7.2 Redox Reactions Mechanism of one electron transfer reaction Outer Mechanism. 7.3 Theories of Redox reactions. 7.4 Potential Energy Diagrams 7.5 Marcus Theory 7.6 Factors Affecting Electron Transfer Reaction Rate 7.7 Inner Sphere Type Reactions 7.8 Let Us Sum Up 7.9 Check Your Progress : They Key 229

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. 230

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 (NH 3 ) 5 X] 2+ into [Cr (H 2 O) 6 ] 2+ or that of [Cr (H 2 O) 5 Cl] 2+ into [Cr (H 2 O) 6 ] 2+. These reactions follow inner sphere or bridge mechanism and can be traced using standard chemical and physical methods 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 Outer Sphere Mechanism In these reactions neither the bonds are formed nor broken up. Hence there is no change in the coordination sphere of metal ions, except their oxidation states are changed. 231

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 [Fe II (CN) 4 ] 4- [Fe III (CN) 6 ] 3-. Which can be studied by labelling of the complexes with a radioactive isotope of Fe or 14 C. [*Fe 2+ (CN) 6 ] 4- + [Fe 3+ (CN) 6 ] 3- [*Fe 3+ (CN) 6 ] 3- + [Fe 2+ (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 10 5 at 25 o C 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. to 232

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 k 2 ) 10-4 for [Co (C 2 O 4 ) 3 ] 3- - [Co(C 2 O 4 ) 3 ] 4- through 10 3 for MnO - 4 MnO 2-4 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 G o 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 Ph 4 As + has little effect, however. These results suggest that a partly desolvated cation accelerates exchange by helping to overcome electrostatic repulsion by formation of a transition 233

234 state such as [Fe (CN) 6 ] 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 MnO MnO - 4 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 (G o must be negative, or the reaction would not take place). This is tantamount to saying that for such reactions the structure of the transition state is more like that of the reactants; hence the activation energy is lowered and the rate is increased. The outer sphere electron transfer may be represented as follows : If oxidant = 0 and reductant = R, then O + R [ O... R ] [ O... R ] [ O... R ]* [ O... R ]* [ O -... R + ] [ O -... R + ] O - + R + First the oxidising agent ( O ) and the reducing agent (R) come closer to form precursor complex. During activation of the precursor complex there is reorganisation of solvent molecules and the bond lengths 234

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 235

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 (H 2 O) 6 ] 3+ + [Fe* (H 2 O) 6 ] 2+ [Fe (H 2 O) 6 ] 2+ + [Fe* (H 2 O) 6 ] 3+ In this reaction also electron transfer takes place through outer sphere or tunneling mechanism. At 25 o 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 *G is. 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, * G ES. Hence the Gibbs energy of complete activation is the sum of all these energies : *G = *G IS + *G OS + *G ES. For electron transfer it is necessary that the energies of the participating electronic orbitals are equal (Frank Condon Principle). In 236

237 this reaction, one electron from t 2g orbital of Fe (II) is transferred to t 2g orbital of Fe (III). The bond-lengths in Fe (II) and Fe (III) complexes are different (In octahedral high spin complex of Fe (II) = 92 pm and that of Fe (III) = 78.5 pm), which indicates the energies of orbitals are not equivalent. If the electron transfer takes place without losing energy, then we will get the product in which bond-length of Fe (II), complex is equal to the characteristics bond length of Fe (III) complex and vice versa. But this will be against the first law of thermodynamics. Actually, it is necessary to give energy for electron transfer. The actual reaction starts with smaller bonds in Fe (II) - b. Complex and larger bonds in Fe (III) complex; unless the energy of participating orbitals become equal. The vibrational stretching and compressions of metal ligand bond help in getting required configuration. 7.4 POTENTIAL ENERGY DIAGRAMS Potential energy diagrams also confirm the relation between molecular motion and electron transfer. For electron transfer it is necessary that there should be coupling in vibrational and electronic motion. The limit of these interactions is related with the energy difference, E, at the crossing of potential energy diagrams (fig. 7.1). If the coupling interaction is strong, distortion of bonds is very less and electron transfer is easy. If the interaction is weak, bond-distortion in high; G will be high and the reaction is slow. As the activated complex rests at the intersection of two curves; while according to the noncrossing rule, molecular potential energy curves of same symmetry states do not cross each other, but divide into upper and lower curves (Fig. 7.1). The indication of non-crossing rule is that, if the reactants distort slowly in their ground states, then they are converted into products, in their ground states itself, following the lowest energy path. 237

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 t 2g orbital, is in non-bonding molecular orbital; and is diffused partially on the ligand due to bonding. During electron 238

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 G o 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 : K 2 = f K 1 K 2 K 239

240 where, K 1 and K 2 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* G 1 + * G 2 + t G o 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* (NH 3 ) 6 ] 3+ + [Co (NH 3 ) 6 ] 2+ [ Co* (NH 3 ) 6 ] 2+ + [Co (NH 3 ) 6 ] 3+ The rate constant of this slow second order reaction is 10-6 M -1 S -1. (i) (ii) (iii) The characteristics of the reactions are as follows : Co N bond Lengths, in Co (II) and Co (III) complexes, are quite different, 2 11 pm and 194 pm respectively, Co (II) complexes are high spin complexes (t 2 g 5 eg 2 ) while Co (III) complexes are low spin complexes (t 2 g 6 eg o ); and after the electron transfer, the configuration in both the complexes probably becomes t 2 g 2 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 ] 4-240

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) (ii) 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). It means, the structure of transition state is analogues to the structure of the reactants, hence activation energy remains low and rate increases. Excited State Outer Sphere Electron. Transfer Reactions Dramatic change are seen in the redox properties of transition metal complexes, when they are in excited states, absorbing energy. In this field, much work has been done with tris (2-2 bipyridine) Ruthenium (II) Cation ([Ru (bpy) 3 ) 2+. Specially, because this has possibility of decomposing water by photochemical reaction. Thus, this can open the path of the production of hydrogen using solar energy. (Creutz and Sution, 1975). hv H 2 O H 2 + ½ O 2 [Ru* (bpy) 3 ]

242 Here ruthenium complex function as photo-sensitizer. G o = 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 V) and with V (-0.86 V V) an outstanding reducing agent (Fig. 7.3) Fig. 7.3 The electronic transfer in the above absorption is a metal to ligand charge transfer, in which one d-electron of ruthenium is exicited and goes to - antibonding orbital of a bipyridine molecule. Hence in the excited state, the structure of [Ru* (bpy) 3 ] + may be written as [Ru(III) (bpy) 2 (bpy)] 2+. The presence of an electron in an antibonding orbital of a ligand makes this excited state cation quite a better reducing agent, as compared the ground state cation. Further the hole thus formed in ruthenium, increases its electron accepting capability, resulting in the fact that this excited state. cation, as compared to its ground state, is also a good oxiding agent. 242

243 7.6 FACTORS AFFECTING ELECTRON TRANSFER REACTION RATE Halpern summarised the factors affecting the rates of direct electron transfer reactions : (i) (ii) (iii) (iv) (v) (vi) Electrostatic repulsion between ions of like charges increases the activation energy; hence the rate of exchange of electrons decreases. When there is no change in the shapes of the molecules, possibility of fast electron transfer always remain. Fast electron transfer takes place, when due to coupling or by anatomic ligand, electrons reach on the surface of the complex. 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. With the increase in the conductivity of the ligand electron transfer increases. 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 G o for the reaction, faster is the reaction. 243

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 (NH 3 ) 5 Cl] 2+ and [Cr (H 2 O) 6 ] 2+ in acidic solution, first investigated by Taube. [Co (NH 3 ) 5 Cl 2+ Cr(H 2 O) 6 2+ [(NH 3 ) 5 Co(III) CI Cr(II) (H 2 O) 5... (7.1) [(NH 3 ) 5 Co Cl Cr (H 2 O) 5 ] 4+ [(NH 3 ) 5 Co] 2+ + [Cr(H 2 O) 5 Cl] (7.2) [(NH) 5 Co] H 2 O [(NH 3 ) 5 Co(H 2 O) (7.3) 244

245 [NH 3 ) 5 Co(H 2 O) H 2 O [Co(H 2 O] NH 3... (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) (ii) Chlorochromium (III) is formed during the reaction. No labeled * Cl atoms are found in the chlorochromium (III) complex when the reaction is carried out in the presence of *Cl - ions, indicating that no ionization of the complexes takes place. If the reaction is carried out in the presence of 36 Cl - in the solution, none of this isotope appears in the chromium (III) complex; this fact provides further support for the bridging mechanism. Because the change in oxidation states of the metal ions is accompanied by transfer of a chlorine atom, the process is often referred to as an atom transfer reaction. Similar reactions occur when the chloride in the above example is replaced by other halide ions, sulphate, phosphate, acetate, succinate, oxalate and maleate. Among halide ions, the effectiveness for bridging purposes (as measured by relative reaction rates) is F - < Cl - < Br - < I -, in accordance with the expected order of ability to transmit an electron and undergo covalent bond-breakage. For the organic ions mentioned, oxalate and maleate (which contain conjugated systems) are considerably more effective than acetate and succinate. 245

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 (NH 3 ) 5 X] 2+ - Cr 2+ (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 Fe 2+ and Fe 3+, 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 FaCl 2+ (aq) and Fe 2+ (aq). The consequences of the inner sphere mechanism of the redox reactions are as follows : (1) Transfer of a ligand from one complex to the other. (2) The rate cannot be faster than the rate of exchange of the ligand in the absence of the redox reaction. (3) The reaction is zero order with respect to one of the complexes and of first order with respect to the other, where bond dissociation takes place. (4) The reaction is first order with respect to the first species if the rate determining step is the attack on the complex. 246

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 Cr 2+ and between CrX 2+ and between Cr 2+ and [Co(NH 3 ) 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(NH 3 ) 5 Cl] 2+ and [Cr (H 2 O) 6 ] 2+ the intermediate dissociate to give...cr (III) and... Co (II). Complexes. (ii) Amongst halide ions, the effectiveness for bridging purposes is in the order...<...<... (iii) When the reaction is carried out in presence of *Cl - ion,... found in the chlorochromium (III) complex, indicating that... of the complex takes place. b (i) Inner sphere mechanism of redox reaction involve transfer of... to the other. (ii) The reaction is... order with respect to one of the complexes and of... takes place. (iii) Electron transfer is rapid only if... system is formed in the intermediate. 247

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. 248

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. 249

250 * Marcus proposed an approximation, giving an equation for rate constant of outer sphere electron transfer reaction. Marcus relation is K 2 = f K 1 K 2 K. where K 1 and K 2 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 G o. * Many redox reactions occur by a ligand briding or inner sphere mechanism e.g. [Co (NH 3 ) 5 Cl] 2+ - [Cr (H 2 O) 6 ] 2+ reaction. The intermediate formed, dissociates to give a 6 coordinated Cr (III) and a 5 coordinated Co (II) complex. * No labelled *Cl atoms are found in the chloro-chromium (III) complex when the reaction is carried out in the presence * Cl ions, indicating inner sphere mechanism and no ionisation of the complex. * Amongst holide ions, the effectiveness for bridging purposes is in the order F - <Cl - < Br - < I - in accordance with the e - transmitting tendency. 250

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 ] (a) (i) Six coordinated Cr (III) and 5 coordinated Co (II) (ii) F - < Cl - < Br - < I - (iii) no lablled *Cl atoms are indicating that no ionisation (b) (i) Ligand from one complex (ii) Zero and of first order where bond dissociation (iii) Conjugate bridge system. 251

252 Unit - 8 METAL-LIGAND EQUILIBRIA IN SOLUTION Structure 8.0 Introduction. 8.1 Objectives. 8.2 Step-wise and Overall Formation Constants Thermodynamic Importance of Stability Constants. 8.3 Factors Affecting Stability Factors related with Metal Factors related with Ligands Chelate effect and its Thermodynamic Origin. 8.4 Methods of Determination of Stability Constants ph-metric method Spectrophotometric method. 8.5 Let Us Sum Up 8.6 Check Your Progress: The Key 252

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 253

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. 254

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 ( ML) M + L ML K 1 = [ M ][ L] ML ML 2 K 2 = ML 2 ML 3 K 3 = ( ML2 ) [ ML][ L] ( ML3 ) [ ML ][ L] 2 Thus ML n-1 + L ML n K n = ( MLn ) [ ML ][ n1 L The equilibrium constants, K 1, K 2, K 3,...K n are called stepwise stability constants. The formation of the complex MLn may also be expressed by the following steps and equilibrium constants. M + L M +2L Thus M + nl B ML, 1 = 2 ( ML) [ M ][ L] ( ML ) [ M ][ L] B 2 ML 2, 2 = 2 n ( MLn) [ M ][ L] B MLn, n = n ]...(8.1) The equilibrium constants, 1, 2, 3,... n are called overall formation or overall stability constants. n is called as n th overall (or cumulative) formation constant or overall stability constants. 255

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 log 10 K 1, log 10 K 2...log 10 K n and log 10 n respectively. Relationship or Interaction Between n and K 1, K 2, K 3,...K n K's and 's are related to one another consider for example, the expression for 3 is:- 3 = ( ML3) [ M ][ L]3 On multiplying both numerator and denominator by [ML] [ML 2 ] and on rearranging we get: 3 = [ ML3 ] [ M ][ L] 3 [ ML][ ML [ ML][ ML 2 2 ] ] = [ ML ] [ ML2 ] [ ML3 ] [ M ][ L] [ ML][ L] [ ML ][ L] 2 = K 1 x K 2 x K 3 Thus n = [ ML ] [ ML2 ] [ M ][ L] [ ML][ L]... [ MLn ] [ ML ][ n1 L ] = K 1 x K 2...K n or n = n 1 K n n n From above relation, it is clear that the overall stability constant n is equal to the product of the successive (i.e. stepwise) stability constants, K 1, K 2, K 3,...K n. This in other words means that the value of stability constants for a given complex is actually made up of a number of stepwise stability constants. 256

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 ( G )changes accompanying complexation. ), entropy ( S ) and free energy 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) negative. The reactions tends to go in the direction written, when G is Enthalpy change ( H ) gives the amount of heat either consumed or liberated per mole of products and is related to the strength of the ligand to metal bonds, compared to that of the metal to solvent bonds. Entropy change ( S ) is related to the change in randomness (the disorder) of a system. As is quite evident from the relation given above (8.2), complex formation is most favoured by the negative enthalpy and positive entropy changes (either of the two or both) as may be expressed by the equation: log K = S H / T R...(8.3) 257

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 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 258

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 259

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 Factors Related With Ligands The properties of the ligands which affect the stability of the metal complexes are as under: 260

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 261

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. 262

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 sp 3 hybridisation in its 4-co-ordinate compounds, rather than with one such as copper which is limited to dsp 2 (planar) hybridisation. Similarly, triethylene tetra amine gives stable complex with metal ions having dsp 2 hybridisation, rather then sp 3 hybridisation Chelate Effect And Its Thermodynamic Origin The chief factor responsible for the stability of the chelate ring is the entropy change which can be viewed statistically or as probability factor. Considering the electronic effect of the donor atom to be the same in the monodentate and the bidentatc ligands, it can be seen that the dissociation of a monodentate from a complex will be higher than that in the chelating bidentate. The dissociation of the M-L bond in monodentate will release the ligand completely from the coordination sphere of the metal, so that it can be easily swept off by the solvent. But the dissociation of one M-L bond for the bidenate ligand does not release the ligand completely (for which simultaneous dissociation at both ends is 263

264 required). Hence the stability constant for metal chelate must be higher. Consider the equilibrium reactions (Fig. 8.4): [Co(NH 3 ) 6 ] en [Co(en) 3 ] NH 3...(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 S o 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 ML n complex is related with the stepwise constants as- n =... (iii) The thermodynamic expression relating equilibrium constant is-... The reaction goes in the direction written when... (iv) CFSE results in the maximum increase in the stability of aquocomplexes of divalent metal ions in the first transition series at d n configuration...and... (v) The Irving Williams order of stability is... (vi) Chelate effect is primarily due to... factor. 264

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, and of entropy respectively. G, H and S are changes of free energy of enthalpy The stepwise or overall stability constant, thermodynamic equilibrium constant gives the value of free energy change, associated 265

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 = [ Mm Ln] m n [ M ] [ L] 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 ph - Metric Method Bjerrum's Method It is a potentiometric method for determining the stability constant for complex formation. Although Bjerrum applied the method primarily to the binding of simple molecules or negative ions to positive metal ions. It may be used with equal success with chelating agents. The theoretical relationship outlined by Bjerrum are not restricted to complex formation but may be applied to any equilibrium process regardless of 266

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 = [HL [L][H ] ] Basic Ligand Acid L + M + KF ML +, KF = [ML ] [L][M ] Basic Ligand metal ion Here Ka and KF are the acid association constant of the ligand and formation constant respectively. Now if CH, Cm and CL are the total amounts in moles/litre of acid (H + ), metal (m + ) and basic ligand (L), we have C H = [H + ] + [HL + ] CL = [L] + [ML + ] + [HL + ] 267

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 + ] = C L - C H + [H + ] - C H [H Ka[H ] ] [M + ] = C m - [ML + ] [L] = C H [H Ka[H ] ] equation in Thus on putting the values of [ML + ], [M + ] and [L] from the above K 1 = [ML ] [M ][L] the value of K 1 can be calculated. For the determination of [ML + ], [M + ] and [L], the values of C H, C L, C m, K a and [H + ], is generally determined potentiometrically using a P H meter. In order to get better results, the ligand must be a medium weak acid or base and the formation constant, K 1, should be within a factor of 10 5 of the value of the acid association constant of the ligand, K a. Irving Rossotti Method Calvin-Bjerrum ph titration technique as adopted by Irving & Rossotti is generally used for determining the proton-ligand and metalligand formation constants. The procedure consists of: (A) Determination of the formation curve of the system. This is expressed as a plot of n (formation function) against pl for metal ligand system and a plot of n A against ph for a proton- 268

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 = Total number of ligand ( L) bound to metal ( M ) Total number of M present in system or n = n io n io.i[mli].[mli]...(8.5) which can be written using equation (8.1) as, n = n io n io.i.β[l].β 1 [L] i 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 io i io.i.β.β H i H i i [H] [L] i [ H o = 1]...(8.7) whereas, pl gives the free ligand exponent and may be defined as. pl = log. 1 [ L] 269

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) (b) (c) Mineral acid Mineral acid + Ligand solution 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 = ( V Y TLo n V ( V TL o )( N E ) 1 V ) 1...(8.8) n = (V iii1 (V V ο n ) (N E ) TL 1 V ) n.tcmο ο o...(8.9) Where V o is the initial volume of the solution, E o, TL o are the initial concentrations of the mineral acid and the reagent respectively and V ', V '' and V ''' are the volume of alkali of a given normality, N, required 270

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 TM o 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 - log 10 n j n o H n 1 ( ) anti log TCL0 n. TCM 0 n 0 iii V V. o V...(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 metalcomplexes are obtained from the formation curve plotted between the values of n and pl - [Fig. 8.2(c)]. The value of formation constants are generally refined using least square method. Fig. 8.2: (a) ph - Titration Curves 271

272 272 (b) Proton-Ligand formation curve (c) Metal-Ligand formation curve 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 = n m x n n m n m m n n m P P C x n m n P n m ] ) ( [ ] ) ( [ 1) ( 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. C 1 = 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 ( P 1)(1 2x) K = 2 C1 [( P 1)( x 1)] Vosburgh and Cooper as well as Katzin and Gebert have extended Job's treatment to systems in which two or more complexes are formed. The ratio of the concentration of metal should not be equal to 1 i.e. nonequimolecular solutions of ligand and metal should be used. 2. Turner Anderson Method Turner and Anderson have modified Job's method and have successfully used for determination of stability constant. By plotting a continuous variation curve for a given range of compositions and then repeating the procedure for more dilute solutions. If the initial concentrations of the metallic ions and ligands are 'a' and 'b' respectively, then K = X ( a x)( b x) where, K = Stability Constant X = Concentration of the complex 273

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 a 1, a 2 and b 1, b 2 represent the concentrations of the metal and the ligand respectively on the two curves, then: K = X ( a x)( b1 1 x ) = X ( a x)( b2 2 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 = ML 2 X ( a x)( b 2x) K = 2 Taking the concentration a 1, a 2 and b 1, b 2 for the same absorbance i.e., the same value of x, we have X ( a x K = 2 1 x)( b1 2 ) X = 2 ( a2 x)( b2 x) The value of x is determined from the relation 4x 2 (a 1 -a 2 +b 1 -b 2 )-x(4a 1 b 1-4a 2 b 2 +b b 2 2 )+ (a 1 b 2 1 -a 2 b 2 2 ) = 0 274

275 i.e. AX 2 + BX + C = 0 Where, A = 4(a 1 -a 2 + b 1 - b 2 ) - B = 4(a 1 b 1-4a 2 b 2 + b b 2 2 ) or b 1 (4a 1 + b 1 ) - b 2 (4a 2 + b 2 ) C = (a 1 b a 2 b 2 2 ) By solving the quadratic equation: K = K = ( B) ( B) ( B 2A ( B 2A 2 2 4AC) 4AC) 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 = M 2 L Taking the concentration a 1, a 2 and b 1, b 2 for the same absorbance i.e., for the same value of X, we have K = X 2 ( a 2x) ( b1 1 x ) = X 2 ( a 2x) ( b2 2 2) or 4 2 (a 1 -a 2 +b 1 -b 2 )-(4a 1 b 1-4a 2 b 2 +a a (a 2 1 b- a 2 2 b 2 ) = 0 i.e. AX 2 + BX + C = 0 Where, A = 4(a 1 -a 2 + b 1 - b 2 ) -B = (4(a 1 b 1-4a 2 b 2 + a a 2 2 ) or b 1 (4a 1 + b 1 ) - b 2 (4a 2 + b 2 ) C = (a 1 b 1 - a 2 1 b 2 ) By solving the quadratic equation, the value of X is determined K = ( B) ( B 2A 2 4AC) or X = ( B) 2 ( B 4AC) 2A Mushran has modified this method so as to suit for 1:3 complexes. 275

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. where c A Aextp [ mx] c is the total analytical concentration (expressed in moles/litre) of the metal or ligand, whichever has the limiting concentration at the point in question. Therefore [MX] = A/A extp. C M = Cm - (mx) = Cm (A/A extp. ) C X = Cm - (mx) = Cx - (A/A extp.) C 276

277 Stability constant K = [ MX ] [ M ][ X ] K = [(Cm A / A extp.) C [(A / A extp.) C] - ][(Cx A / A extp.)c ] Where A = Absorbance at the metal ligand ratio. A extp. = The extrapolated value of Absorbance. Cm = Concentration of the metal at equivalence point. Cx = Concentration of the ligand at equivalence point. C = Total analytical concentration of the ligand. When metal ligand ratio and the ratio shown by extrapolation do not be on the same ordinate, then the value Cx and C will not be the same. C is calculated at the point of intersection of the extrapolated curve. 4. Raghav Rao's Method Subbarama Rao and Raghav Rao used job's method of continuous variation and molar ratio method for determination of stability constants. They used equimolar solutions of metal and ligand with optical density as he index property. This method is also known as graphical method. Reddy and Seshoish used the same graphical method using conductance and optical density as the index property. 277

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 ML n 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 K 1, K 2, K 3,...K n 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 ML n, 278

279 as : n = K 1, K 2, K 3,...K n or n n n K n n1 The equilibrium constant is related to the thermodynamic expression as follows: - RT log K = G = H - T S. The factors affecting stability of complexes are mainly related with the metal ion and the ligands. The factors due to metal are primarily related with the size of the ion, its charge, possibility of -bonding and CFSE gained. Stability is proportional with the charge and ionic potential (e/rratio) but is inversely proportional with the size of the metal ion. 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. 279

280 Shape of the ligand molecule also affects stability e.g. while triethylene teramine gives complex with metal ions having dsp 2 hybridisation(sq. planar geometry);, I, II triamminotriethylamine gives stable complex with metal ions having sp 3 hybridisation (Tetrahedral geometry). Chelates are more stable compared to non chelates. Stability increases with the number of rings formed per mole of the ligand e.g. M(en) < M(trien) < M(EDTA) (1 ring) (2 ring) (5 ring) Higher stability of chelates is mainly related with the entropy factor. The stability constants of metal complexes in solution are determined generally using two methods: one the potentiometer (ph) titration method due to Bjerrum and its modification by Irving and Rossotti; and the other one spectrophotometer methods due to job and its modification by Turner-Anderson. In Irving Rossotti method stability constants are computed by plotting formation curves, between n (the formation function) and pl -. n is the average number of ligand bound per metal atom or ion; while pl is the free ligand exponent; log I [L] According to half integral method: The value of pl - at 0.5 n = Log K 1 The value of pl - at 1.5 n = Log K 2 and so on. The values of stability constants are generally refined by least square method. 280

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 βn n n n K n 1 (iii) - RT log K = G = H - T S. (iv) (v) G is negative. 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 - I = [L] (iv) Between n and PL - (v) (vi) Job's method used for determination of stability constants by plotting continuous variation curve...of composition. The total absorbance of. 281

282 Unit - 9 METAL CLUSTERS Structure 9.0 Introduction. 9.1 Objectives. 9.2 Boranes and Higher Boranes Wade's Rule Closo-Boranes Nido-Boranes Arachno-Boranes Structural Interrelation Synthesis Reactions. 9.3 Carboranes Synthesis Properties Structures. 9.4 Metalloboranes and Metallocarboranes Properties. 9.5 Metal Carbonyl Halides. 9.6 Compounds with metal-metal multiple bonds. 9.7 Let Us Sum Up. 9.8 Check Your Progress: The Key. 282

283 9.0 INTRODUCTION Closed polyhedrons play important part in the synthesis of clustermolecules in inorganic chemistry. These cluster-molecules include polyhedral boranes, carboranes and metalloboranes and metallo carboranes; organometallic clusters and metal halide clusters. The definition of metal clusters includes those molecular complexes in which metal-metal bonds form a triangular or a large closed structure. This definition does not include linear M-M-M bonded compounds or those cage like structures in which metal atoms, in closed structures are interlinked through ligands, forming M-L-M bonds. Presence of metal-metal (M-M) bond in these molecules may be ascertained with the help of data of bond lengths and also the stability of compounds. As amongst d-block groups, metal-metal bond strength gradually increases moving down a group, hence d-block metal in fourth and fifth periods of the periodic table form M-M bonded compounds in large number. 9.1 OBJECTIVES The main aim of this unit is to study the nature, methods of preparation and structures of metal-clusters. After going through this unit you should be able to: describe boranes and higher boranes with reference to their classification, synthesis reactions and structures; discuss carboranes and explain their synthesis and properties in the light of their structures; describe metalloboranes and metallocarboranes in relation with carboranes; explain structures of metal carbonyl halides; and 283

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 BH 3, but this compound is unstable at the room temperature. However, higher hydrides like B 2 H 6 (diborance). B 4 H 12 (tetraborane), B 6 H 10 (hexaborane), B 10 H 14 (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 He also proposed the concept of completely delocalised molecular orbitals to explain structures of boron polyhedrons. He established icosahedral structure of [B 12 H 12 ] Fig 9.2. Fig. 9.1: 3C, 2e bond in B 2 H 6 284

285 Fig. 9.2: B 12 H 12 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 sp 3 hybridisation are placed at the corners of a equilateral triangle (Fig. 9.3). Fig. 9.3: B-B-B bond Wade's Rule In 1970 K. Wade gave a rule relating the number of electrons in the higher borane molecules with their formulae and shapes. Using these rules one can predict the general shapes of the molecules from their formulae. These rules are also applicable on carboranes and other polyhderal molecules called 'Deltahedral's Deltahedrons are so called, as they are composed of delta,, shaped triangular faces. 285

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 B 4 H 10, four BH units contribute 8 electrons (4x2 = 8) and six extra hydrogens give six electrons thus B 4 H 10 has total 14 skeletal electrons Fig 9.4 gives the molecular energy diagram of [B 6 H 6 ] 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 (a 1g, t 1u and t 2g ). t ui e g t 2g t 2u t 2g t 1u a 1g Fig. 9.4: Skeletal molecular energy diagram of [B 6 H 6 ] 2-286

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 [B n H n ] 2- n+1 [B 5 H 5 ] 2- to [B 12 H 12 ] 2- Nido [B n H n+4 ] n+2 B 2 H 6, B 5 H 9, B 6 H 19 Arachno [B n H n+6 ] n+3 B 4 H 10, B 5 H 11 Hypo [B n H n+8 ] n+4 Only derivatives are known Closo Boranes These are closed structured (Closo, Greak, meaning cage) boranes with the molecular formula [B n H n ] 2- and skeletal electrons = n+1 pairs (= 2n+2 electrons). In this structure, there is one boron atom placed at each apex and there are no B-H-B bonds present in the molecule. All the member of the series from n=5 to 12 are known. [B 5 H 5 ] 2- is trigonal bipyramidal, [B 6 H 6 ] 2- is octahedral and [B 12 H 12 ] 2- is icosahedral. All are stable on heating and are quite inert Nido-Boranes These boranes have nest (Nido, Latin, meaning Nest) like structure. Their general formula is B n H n +4 and have (n+2) pairs = 2n+4 skeletal electrons on removing one boron atom from an apex of closo structure, nido structure is obtained. Because, of the lost boron atom, these boranes have extra hydrogens for completing the valency. The polyhedra in this series have B-H-B bridge bonds in addition to B-B bonds. They are 287

288 comparatively less stable than 'Closo', but more than 'Arachno' on heating Arachno-Boranes These boranes have the general formula (B n H n +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 Structural Inter-relation The structural interrelation between closo, nido arachno species is shown in Fig Arachno B 4 H 10 Fig. 9.5: 288

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. [B 6 H 6 ] 2- ion and adding four hydrogens, we get square pyramidal nido- B 5 H 9 borane. On repeating same process on nido B 5 H 9 (i.e. removing one BH unit and adding two hydrogen's), we get butterfly shaped arachno. B 4 H 10. 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) Synthesis The simplest method for synthesis of higher boranes is the controlled pyrolysis of diborance, B 2 H 6 it is a gas phase reaction, BH 3 formed in the first step reacts with borane to give higher boranes: B 2 H 6(g) 2BH 3(g) B 2 H 6(g) + BH 3(g) B 3 H 7(g) + H 2(g) B 3 H 7(g) + BH 3(g) B 4 H 10(g) B 2 H 6(g) + BH 3(g) B 3 H 9(g) [B 3 H 8 ] - (g) + H + 5[B 3 H 6 ] (g) [B 12 H 12 ] 2- (g) + 3[BH 4 ] - (g) + 8H 2(g) 2[BH 4 ] (g) + 5B 2 H 6(g) [B 12 H 12 ] H 2 289

290 Closp Nido Arachno Fig. 9.6: Interrelation between closo, nido and arachno-boranes Reactions The important reactions of higher boranes are with Lewis bases, which involve removal of BH 2 or BH n from the cluster, growth of the cluster or removal of one or more number of protons: 1. Decomposition by Lewis-bases: B 4 H NH 3 [BH 2 (NH 3 ) 2 ] + [B 3 H 8 ] The reaction is analogous to the reaction of diborane with ammonia. 290

291 2. Deprotonation : Higher boranes give deprotonation reaction easily rather than decomposition: B 4 H 10 + N(CH 3 ) 3 [HN(NH 3 ) 3 ] + [B 10 H 13 ] - This deprotonation takes place from 3c, 2e BHB-bond. The bronsted acidity of boranes increases with their size: required: B 4 H 10 < B 5 H 9 < B 10 H 14 For deprotonation of B 5 H 9 strong-base like Li 4 (CH 3 ) 4 is B 5 H 9 + Li(CH 3 ) Li + [B 5 H 8 ] - + CH 4 3. Cluster Building: Reactions of borane with borohydride are important with respect to synthesis of higher boranes: 5K[B 9 H 14 ] + 2 B 5 H 9 5K[B 11 H 14 ] + 9 H 2 4. Electrophilic displacement of proton: Electrophilic displacement of proton by the catalytic activity of Lewis acids like AlCl 3 is the basis of alkylation and halogenation of boranes: B 5 H 9 + CH 3 Cl AlCl [CH 3 B 5 H 8 ] + HCl 3 291

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 B n H -2 n anions with B - replaced by isoelectronic carbon. These carboranes have the general formula. C 2 B n+2 (n=3) to 12. The important member is C 2 B 10 H 12 (Fig. 9.7). Which is isoelectronic with [B 12 H 12 ] 2- similarly B 4 C 2 H 6 is isoelectronic with [B 6 H 6 ] 2-. (A) 31, 2, C 2 B 10 H 12 (B) C 2 B 4 H 6 Fig Nido Carboranes: They are having an open case structure in which some framework members are attached likely by hydrogen bridges. These are derived formally from one or other of several borones. These contain one to four carbon atoms in the skeleton. In addition to the above types of carboranes, there are a number of carboranes with an additional heteroatom such as phosphorus built into the basic structure and a family of metallo carboranes, some of which are similar to ferrocene. One peculiar feature common to all carboranes is that to date no compound has 293

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 CB 5 H 9 is name as monocarbonido hexaborane (9). Similarly, the three isomers of C 2 B 10 H 12 are named as 1, 2; 1, 7 and 1, 12 dicarbocloso-dodecaborane (12). ii. Number of atoms in these structure are counted by starting the numbering from that in the apical position and proceeding through successive rings in a clockwise direction. This rule is important in naming the isomers. Closo-Carboranes or Closed Cage Carboranes These carboranes are having general formula C 2 B n H n+2 (n=3 to 10) in which the constituents are only terminal. These are isoelectronic with the corresponding [B n H n ] 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 C 2 B n skeleton. They are considered in three groups. a. small, n = 3-5 b. large, n = 6-10 and 294

295 9.3.2 Preparation: c. dicarbo-closo-dodecaborone I(a) The Small Closo Carboranes (C 2 B n H n+2 where n = 3 to 5) B 5 H 9 + C 2 H 2 o C 490 1,5 - C 2 B 3 H 5 + 1,6 - C 2 B 4 H 6 +2,4 - C 2 B 5 H 7 Example - The closo hexaborane isomers, C 2 B n H 6, (b) The Large Closo Carboranes (C 2 B 2 H n+2 where n = 6 to 9) The first three members of this group of carboranes are obtained by the thermolysis of 1,3 - C 2 B 7 H 13 and 1,3 - C 2 B 2 H 12. Example : C 2 B 6 H 8 is made from hexaborane (10) and dimethylacetylene. The structure of 1,7 - Me 2 C 2 B 6 H 6 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 B 10 H 12 L 2. B 10 H L 2 H R B 10 H 12 L 2 2 C 2 R 2 L 2 B 10 H 10 + H 2 + 2L Example: C 2 B 10 H 12 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: CB 5 H 9, C 2 B 4 H 8, C 3 B 3 H 7, C 4 B 2 H 6 etc. Preparation: The smaller nido-carboranes are generally prepared by reacting a borane with acetylene under mild conditions. 295

296 Example: B 5 H 9 and C 2 H 2 undergo reaction in the gas phase at 215 o C to give mainly the nidocarborane 2,3 - C 2 B 4 H 8 together with methyl derivatives of CB 5 H 9. 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) CB 5 H 7 is formed by passing silent electric discharge through 1-methyl pentaborane (9). ii. The only example isoelectronic with B 5 H 9 is 1,2- dicarbonido - pentaborane(7), C 2 B 3 H 7, which is prepared as follows: B 4 H 10 + C 2 H 2 50 o C C 2 B 3 H 7 (3-4 % yield) iii. Monocarbonidohexaborane (9), CB 5 H 9 is formed from ethyldifluoroborane and lithium. The nido-carboranes are formally related to B 6 H 10. All are having eight pairs of electrons which are bonding the six cage atoms together. Large Nido-Carborane: Dicarbo-nido-undecaborane, C 2 B 9 H 13, is the second member of the class of nido-carboranes C 2 B n H n+4 (n =4 or 9),. The parent carborane and its substituted derivatives can be prepared by the base degradation of ortho-carborane (1,2-dicarboclosododecaborane (C 2 B 10 H 12 ). is formed. MeO 1,2 - C 2 H 10 H 12 H C 2 B 9 H 12 1 C 2 B 9 H 13 When C 2 B 9 H 13 is heated, the closo-undeca-borone (11) cage 296

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 500 o C, it is converted into 1, 7 isomer i.e. meta or neo isomer; while at 700 o C it is concerted to 1, 12 isomer i.e. para-isomer (Fig. 9.8) Fig. 9.5: (a) C 2 B 10 H 12 (b) 1,7 C 2 B 10 H 12 (c) 1,12 C 2 B 10 H 12 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 C 2 B 10 H 12, which are described below. Various substitution reactions have been studied and the hydrogen atoms bonded to carbon are weakly acidic. All three of the icosahedral isomers are stable both to heat and to chemical attack, and much more so than decaborane (14). They are white crystalline solids which resist both strong oxidizing agents and strong reducing agents and are also stable to hydrolysis. This is important because it allows reactions to be carried out on substitutions, often under quite drastic conditions, without destroying the cage structure, rather as 297

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-C 2 B 4 H 8 is converted to the closocarboranes C 2 B 3 H 5, C 2 B 4 H 6 and C 2 B 5 H 7 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, C 2 B 4 H 8 + NaH diglyme Na + C 2 B 4 H H 2 Similarly with LiC 4 H 9, Lithium derivative is former: B 10 C 2 H LiC 4 H 9 B 10 C 2 H 10 Li 2 + 2C 4 H 10 The Sodium derivative with FeCl 3 gives Fe-derivative: 2Na 2 [B 9 C 2 H 11 ] + FeCl 3 2NaCl + Na 2 [Fe(B 9 B 2 C 11 ) 2 ] 298

299 9.3.3 Structures Structural studies of carboranes have been done using X-ray analysis and nmr studies. The C 2 B 3, C 2 B 4 and C 2 B 5 closo-carboranes, for example, have trigonal bipyramidal, octahedral and pentagonal bipyramidal skeletal structrues respectively, and positional isomers have been identified. The icosahedra structure is similar to that of B 12 H 2-12 (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 C 3 B 3 H 7 is shown in Fig 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 BH 2 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, C 2 B 3 H 7 has a square pyramidal structure that is formally derived from that of B 5 H 9, with two BH 2 replaced by 2CH. Fig

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 [B 11 H 11 AlCH 3 ] 2- (Fig. 9.10). It is prepared by the action of trim ethyl aluminium [Al(CH 3 ) 3 ] 2 with Na 2 [B 11 H 13 ]: Al 2 (CH 3 ) 6 + 2[B 11 H 13 ] 2-2[B 11 H 11 AlCH 3 ] CH 4 Fig. 9.10: Closo [B 11 H 11 AlCH 3 ] 2- The hydrogen attached with carbon in closo- B 10 C 2 H 12 is slightly acidic. This can be substituted by butyl lithium or Grignard's reagent to get lithium or magnesium metallocarboranes: 2C 4 H 9 Li + C 2 H 2 B 10 H 10 C 2 Li 2 B 10 H 10 2RMgBr + C 2 H 2 B 10 H 10 [CMgBr] 2 B 10 H R-H Similarly, [C 2 B 9 H 11 ] 2- ion, reacts with FeCl 2, BrRe(CO) 5 or BrMn(CO) 5 to give Fe, Re or Mn derivatives: 300

301 2[C 2 B 9 H 11 ] 2- + FeCl 2 [(C 2 B 9 H 11 ) 2 Fe] Cl - [C 2 B 9 H 11 ] 2- + BrRe(CO) 5 [C 2 B 9 H 11.Re(CO) 3 ] - + Br - + 2CO [C 2 BgH 11 ] 2- + BrMn(CO) 5 [C 2 B 9 H 11.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 2- of the type (C 2 B 9 H 11 )M(CO) 3 (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 2- carborane structures is not restricted to the C 2 B 9 H 11 ion; there are a 3- number formed on the same principle by CB 10 H 11 and some of its amine-substituted derivatives (e.g. [(CB 10 H 11 ) 2 Cr] 3- and C 2 B 4 H 3-6 ) also give complexes, and it may be noted, some of these are nido-anions. Thus [1,6 C 2 B 7 H 9 ) 2 Co] - has the structure shown below (Fig. 9.11), the ion being derived from 1,3-C 2 B 7 H 13. (a) (b) (c) Fig. 9.11: (a) Carbonyl metallocene (b) Carbonyl Cyclopentadieny (c) Carbolyl Carbonyl Compound 301

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 B 3 H 7 [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) Properties Fig. 9.12: Structure of [Fe(CO 3 )B 4 H 8 ] 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- PCl 3 + C 2 PhL 2 B 10 H 10 (C 2 PhB 10 H 10 ) 2 Pl Ph 3 PAuCl + C 2 RLiB 10 H 10 Ph 3 AuC(Cr)B 10 H 10l 2(C 6 H 5 )2PCl + C 2 Li 2 B 10 H 10 (C 6 H 5 ) 2 PC-CP(C 6 H 5 ) 2 N ( CO ) 4 i l l - B 10 H 10 - OC CO Ni (C 6 H 5 ) 2 PC CP (C 6 H 5 ) 2 B 10 H

303 Similarly, derivatives of mercury and other metals( -bonded) have also been obtained, RC 2 LiB 10 H Cl H B 10 H 10 RC 2 HgC 2 RB 10 H 10 Ph 3 PAuCl + C 2 RLiB 10 H 10 Ph 3 PAuC(CR)B 10 H 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 Rh 6 (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, Rh 6 (CO) 16 should have closo- structure Some examples showing inter-relation between cluster-valency electrons and structures are given in Table

304 No. of Metal Atoms Geometry Metal Skeleton Structure Table 9.2 Bonding Molecular Orbital No. of Cluster electron Examples 1. Monomer 9 18 Ni(CO) 4 2. Dimer Fe(CO) 9, Mn 2 (CO) Triangle Os 3 (CO) 12, Co 3 (CO) 9 CH 4. Tetrahedron Co 4 (CO) 12, Rh 4 (CO) 12 Butterfly Re 4 (CO) 16 2-, [Fe 4 (CO) 12 C ] 2- Square Os 4 (CO) 16, Pt 4 (O 2 CMe) 8 5. TBP Os 5 (CO) 16 Octahedral Fe 5 (CO) 15 C 6. Trigonal prism Ru 6 (CO) 17 C [Rh 6 (CO) 15 C] 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). 304

305 Tetrahedron Butterfly Square Planar Fig Synthesis: 1. Pyrolytic Synthesis: 2CO 2 (CO) 8 CO 4 (CO) CO 2. Redox Condensation: Ni(CO) 4 + [Ni 5 (CO) 12 ] 2- [Ni 6 (CO) 12 ] CO 3. Ston's Method: Condensation of meatl carbonyls with unsaturated metal carbonyls: (CO) 5 Mo = C(OMe)Ph + Pt(Cod) 2 (CO) 5 Mo.Pt(Cod)(OMe)Ph Cp.W(CO) 2 (C.tol) + Co 2 (CO) 8 (Cp)(CO) 2 W.Co 2 (CO) 6 C.tol Reactions: 1. Substitution Fragmentation: Fe 3 (CO) 12 + P.Ph 3 Fe 3 (CO) 11 (Ph 3 ) + Fe 3 (CO) 10 (P.Ph 3 ) 2 + Fe(CO) 5 + Fe(CO) 4 (P.Ph 3 ) + Fe(CO) 3 (PPh 3 ) 2 + CO 2. Prolongation: [Fe 3 (CO) 11 ] 2- + H + [Fe 3 H(CO) 11 ] - 3. Cluster Catalytic Ligand Transformation: eg. in Os cluster. 305

306 9.5 Metal Halide clusters Although the first information of metal halide clusters was given in 12 th Century in the form of calomel, but dimeric nature of mercurous ion could be established in 20 th Century only. But now number of metal halide clusters are known. Dinuclear Complexes: Most important dinuclear species is [Re 2 X 8 ] 2- (Fig. 9.14). Fig. 9.14: Structure of [Re 2 Cl 8 ] 2- Analogous to [Re 2 X 8 ] 2- ion, in which very small M-M distance and eclipsed configuration of chlorine atoms are present, is [Mo 2 Cl 8 ] 2- and [W 2 Cl 9 ] 3- (Fig. 9.15). Fig. 9.15: Structure of [W 2 Cl 9 ] 3-306

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, [ReCl 3 ] 3 or Fe 3 Cl 9 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 Fig. 9.16: Structure of [W 2 Cl 9 ] 3- Tetra nuclear Clusters: Only a few examples of tetranuclear clusters of halides and oxides are known. Most important example is the dimeric [Mo 2 Cl 8 ] 4- cluster giving a tetra nuclear molecule : 307

308 Hexanuclear clusters: Hexanuclear Clusters or Mo, Nb and Ta halides are well known. Two species are known, one with molecular formula M 6 X 12 or [M 6 X 8 ]X 4 and the other with molecular formula [M 6 X 14 ] 2-. Molybdenum forms cluster of the type [M 6 X 8 ]X 4. [M 6 X 8 ] 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 [M 6 Cl 8 ] 4+ Niobium and tantalum give clusters of M 6 Cl 12 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 [M 6 X 12 ] Similarly, Nb, Ta and Zr give clusters of [Nb 6 Cl 12 L 6 ] 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 [M 6 X 18 ] 2+ type clusters (Fig. 9.19): 308

309 Fig. 9.19: Structure of [M 6 X 18 ] 2+ (Fig. 9.20). Solid [Mo 6 Cl 14 ] 2- species is derived from MoCl 2 as its hexamer Fig COMPOUNDS WITH METAL-METAL MULTIPLE BONDS As has been shown earlier, the earlier metals in d-block series in their lower oxidation states have tendency to form metal-metal multiple bonds. These metal-metal bonds may be present in smaller molecules and also in macro-chain solids. 309

310 Chevrel-phases : Chevrel phases generally involve tertiary molybdenum chalcogenides, MxMoX 6, 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, PbMo 6 S 8. 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 Mo 6 S 8 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 = Mo, o = s, 0 = Pb Zintle anions and cations: In 19 th 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 Sn 4-9, Pb 4-7, Pb 4-9,Sb 3-7 and Bi 3-3 were discovered. In 1975 cryptate salts of these anions were also obtained. Some cations, Bi 5+ 9, Te 4+ 6, etc were also prepared 310