Hydrogen Bonds Involving Transition Metal Centers Acting As Proton

Research: Science and Education

Hydrogen Bonds Involving Transition Metal Centers Acting As Proton Acceptors Antonio Martín* Departamento de Química Inorgánica, Instituto de Ciencia de Materiales de Aragón, Universidad de Zaragoza-C.S.I.C. E-50009 Zaragoza, Spain

Since Huggins, Latimer, and Rodebush first described the hydrogen bond in 1920 (1), this kind of interaction has received a great deal of attention (2). After ionic and covalent bonds, hydrogen bonds are the most common and strongest interatomic interaction. They are of capital importance in many areas of chemistry, physics, biology, crystallography, mineralogy, and other related areas. Background Hydrogen bonding can be defined as a weak, secondary interaction between a lone electron pair and a hydrogen atom bound to an electronegative residue. The hydrogen atom has an electropositive character and acts as a Lewis acid, whereas the Y atom behaves as a Lewis base. δ{

δ+

X — H ???? : Y The X–H fragment is usually referred to as the proton donor and Y as the proton acceptor. The most common proton donors are N–H and O–H groups, but other examples are also known involving P–H, S–H, F–H, Cl–H, and Br–H moieties as proton donors. C–H can also be a proton donor under certain circumstances: the C–H bond has to have some polarity and thus C–H????Y hydrogen bonds have been found when the carbon atom is bound to electronegative groups (as in CHCl3) or is in a hybridized sp state. The proton acceptor has to have lone electron pairs; that is, it must be a Lewis base. Among the “classic” proton acceptors are N, O, halogen, S, and P atoms. Other less typical proton acceptors are Se atoms (3) or C atoms of phosphonium ylide groups (Ph 3P +–{CR2) in which the C atom has a negative charge (4 ). The π electron density of unsaturated or aromatic systems can also act as a Lewis base (5). These are the only cases in which the C atoms behave as proton acceptors. Hydride is another fairly striking proton acceptor, giving X–H????H–M hydrogen bond systems (6 ). Several theoretical models have been used to explain the hydrogen bond. The simplest and most commonly used is the electrostatic model, based on ion–dipole or dipole–dipole interactions, as is supported by the observed preference of hydrogen bonds for cases involving electronegative atoms. However, this simple model has some deficiencies. Thus, for example, the shortness of some hydrogen bonds might be expected to give rise to repulsive interactions, which could in turn weaken the hydrogen bond, but this is not observed. In a molecular orbital approach there is a covalent interaction spread over the three X–Y????H atoms. This model is sometimes called the 3-center–4-electron (3c-4e) model, the three centers being the X, H, and Y atoms and the four electrons the two from the X–H bond plus the two donated by the Y atom. *Email: [email protected].

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Hydrogen bonding is a directional interaction. This means that in the vast majority of the X–H????Y systems the values of the X–H–Y angles are clustered around 160–180°. The H????Y distance is related to the strength of the interaction. For weak interactions such as C–H????(N,O), the usual H????(N,O) distance ranges between 2.20 and 3.00 Å. For stronger interactions, shorter H????Y distances have been found (1.80–2.00 Å for N–H????O and 1.60–1.80 Å for O–H????O systems). Even shorter distances can be found with more electronegative atoms (5), the extreme case being that of some “symmetric” O????H????O or F????H????F systems as found in KHF2, with H????F distances of 1.15 Å. The hydrogen bond energy is around 20–40 kJ mol{1 for strong interactions (for example 29 kJ mol{1 for F–H????F in gaseous HF or 21 kJ mol{1 for O–H????O in solid water) and 5–20 kJ mol{1 for weaker C–H????Y hydrogen bonds. These are low values when compared with the energy of a covalent bond (e.g., 347 kJ mol{1 for C–C). Apart from these more usual hydrogen bonds there are some multicenter or bifurcated bonds (Fig. 1) (2b). Of these, only I is found in remarkable numbers (≈20% when X = N and Y = O= C) (8), the rest being very rare. Y O

Y

H H

C X

Y

H

O

C O

Y'

I

III Y

Y H

H Y'

Y" H

H

II

O

Y'

Y' Y"

O H

H

Y'''

Y"

IV

V

Figure 1. Multicenter hydrogen bonds.

The existence of hydrogen bonds has several spectroscopic consequences. They can be detected by infrared and Raman spectroscopy mainly because of a decrease in the νXH due to the weakening of the X–H bond and the broadening of the band. In the proton nuclear magnetic resonance spectra, the signal corresponding to the hydrogen involved in the interaction is shifted to a higher frequency. Despite their low bond energy the importance of hydrogen bonds in many aspects of chemistry must not be underestimated. Intermolecular hydrogen bonding is responsible for some of the physical properties of compounds of crucial importance, such as water. If one studies the boiling points of group 16 hydrides it can be found that by extrapolating

Journal of Chemical Education • Vol. 76 No. 4 April 1999 • JChemEd.chem.wisc.edu


the values for H2Te, H2Se, and H2S, the boiling point of H 2O should be around {100 °C, whereas it is actually +100 °C. The reason is that in liquid water the molecules are strongly associated by hydrogen bonds. Similar “abnormally high” boiling points compared to their group congeners are found for NH3 and HF, in which there are hydrogen bond interactions. For HF in the solid and liquid state, the H atom of one molecule is linked to the F atom of another molecule forming an infinite zigzag chain (Fig. 2a). Some hydrogen bonds even persist in the gas phase. In group 14, however, the boiling point of CH4 follows the trend set by SiH4, GeH4, and SnH4 owing to the inexistence of hydrogen bonding. Similar patterns can be seen in the melting points and enthalpies of vaporization of the hydrides, indicating hydrogen bonding in NH3, H2O, and HF, but not in CH4. O

O H H

H F

F

F

B F

H

C H

H

O

H

O B

O

C O

O

H

H

O

H

B O

H

(a) H

H

H O

O H

O

O O

F

H

H

O

H

F

B

O

O H

H

B

O O

Hydrogen bonding, both intermolecular and intramolecular, has a special significance in biological systems (11). It is responsible, for example, for the linking of polypeptide chains in proteins and of base pairs in nucleic acids (12). Because hydrogen bonds are directional they are sensitive to stereochemistry and, owing to their low binding energy, they can be switched on and off with energies that are within the range of thermal fluctuations at physiological temperatures. Both factors are very important, for example, in the processes of enzyme recognition of substrates and in the formation of enzyme–substrate complexes, which is the first step in the catalytic action of enzymes. The structural complementarity between the substrate and the active site of the enzyme is the basis of substrate specificity. Successful binding of a substrate to an enzyme occurs when the shape of the active-site cleft fits the subsite molecular structure. X-ray crystallographic studies of enzymes have, in fact, demonstrated that substrates fit snugly into the active-site cavity on the surface, resulting in complexes of a well-defined structure with many contacts between the enzymes and their substrates. Some of these contacts are hydrogen bonds, and although each on its own is weak, because they occur at numerous sites the resultant interactive force between substrate and enzyme is considerable (13). Transition Metals as Proton Acceptors

O B

H

H

O

(b)

O

(c)

Figure 2. Intermolecular hydrogen bonds. (a) Zigzag chain in solid HF; (b) formic acid dimer; (c) fragment of a layer of crystalline H3BO3.

The hydrogen bonds present in NH3, H2O, and HF are intermolecular, as are those found in formic acid and boric acid (Fig. 2). Intermolecular hydrogen bonding is the most important directional interaction in supramolecular chemistry (9). The majority of well-known and structurally robust intermolecular assemblies are based on hydrogen bonding. There are good examples of systematic crystal engineering using the directional properties of strong hydrogen bonds (9, 10). In some cases the hydrogen atom interacts with a proton acceptor located in the same molecule. These are intramolecular hydrogen bonds and they can also substantially affect some properties of the compounds. Thus, in o-nitrophenol there is a hydrogen bond between the OH hydrogen and one of the oxygen atoms of the nitro group (see Fig. 3a), giving rise to a decrease in its acidity compared with its meta and para isomers. Hydrogen bonding has the opposite effect in salicylic acid (o-hydroxybenzoic acid, Fig. 3b). The hydrogen bond between the hydrogen of the 2-hydroxy substituent and the hydroxy oxygen atom of the carboxylic group enhances the acidity of the carboxylic hydrogen. Consequently, it has been observed that salicylic acid is a much stronger acid than its meta and para analogs.

In the last few years an increasing number of hydrogenbonding interactions in which the proton acceptor is a metal center of a coordination complex, M????H–X, have been reported in the literature (14 ). These interactions are substantially different from the well known “agostic” interactions (15). In the latter, the metal center acts as a Lewis acid, generally receiving electron density from a C–H bond. This is a 3-center–2-electron (3c-2e) bond system. The three centers are the metal and the C and H atoms, and the two electrons are those in the C–H bond. An empty orbital of the metal is involved to house the donated electron density. This is an electron-deficient bond system of the same nature as that which is present in the diborane molecule, B2H6. On the other hand, the M????H–X hydrogen bonding systems are substantially similar to the “classic” hydrogen bonds. The metal atom is the Lewis base that has a filled orbital with an electron pair that can interact with an electropositive hydrogen atom, or, using a molecular orbital method language, the electron pair is donated to create a 3-center–4-electron (3c-4e) system. These hydrogen bonds therefore are favored by electron rich metals such as late transition metals, especially in low oxidation states. In fact, all the M????H–X hydrogen bonds described so far involve metal centers with d8 or d10 electron configuration. In the case of the square-planar Pt(II) (d8) complexes, the highest-energy d orbital is the 5dx {y , which is unoccupied, and X the other four filled d orbitals, 5dxy , 5dyz, H 5dxz , and 5dz , have very similar energy values, their exact order being dependent on the nature of the ligands. The L orbital with a symmetry that is suitable L L for interaction with hydrogen can be easily visualized as the 5dz (see Fig. 4). L The differences between agostic and hydrogen M????H bonding interactions Figure 4. The 5dz gives rise to some different structural and orbital. 2

2

2

O

O

N

C O

O

H

O O

H

H

2

(a)

(b)

Figure 3. Intramolecular hydrogen bonds. (a) o-Nitrophenol; (b) salicylic acid.

JChemEd.chem.wisc.edu • Vol. 76 No. 4 April 1999 • Journal of Chemical Education

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spectroscopic features, which will be compared and discussed later (see Table 1). Both intermolecular and intramolecular M????H–X hydrogen bonds have been described, the latter being more common. Both types have been studied using different structural and spectroscopic techniques both in the solid state and in solution. The most important techniques for studying hydrogen bonds in the solid state are X-ray and neutron diffraction. Of these, the more powerful is neutron diffraction because it is able to precisely locate the position of the hydrogen atoms, whereas X-ray diffraction can seldom provide this information. Nevertheless, a careful X-ray diffraction experiment can give a good estimate of the hydrogen atom’s coordinates. Furthermore, X-ray diffraction is a much more accessible technique than neutron diffraction and there is, therefore, much more X-ray data available. It should be emphasized that the positions of H atoms as determined by X-ray diffraction are different from those determined by neutron diffraction. The former shows the H atom apparently closer to the heavier atom to which it is bonded. The typical difference in the O–H distance is about 0.15 Å (neutron diffraction, 1.0 Å; X-ray, 0.85 Å). This effect appears to arise from the asphericity of the electron distribution due to chemical bonding. This affects the X-ray scattering factor of the atom, which depends on the orbital electrons but not the neutron scattering factor (which, for a diamagnetic atom, is purely nuclear). Methods for refining crystal structures involve the use of calculated atomic scattering factors so that, if a spherical electron distribution around an atom is assumed and the position of an atomic nucleus is determined as the center of gravity of its electron cloud, the X-ray position may be different from that determined by neutron diffraction (16 ). In solution, the most useful technique for studying M????H–X hydrogen bonds is nuclear magnetic resonance (NMR) spectroscopy, as described below. Infrared spectroscopy has also been used in some cases. Since hydrogen bonds are weak, it may be possible that interactions detected in the solid state are not present in solution.

M????H–X Intramolecular Hydrogen Bonds To date, most of the M????H–X hydrogen bonds reported in the scientific literature are intramolecular. They involve hydrogen atoms belonging to a fragment of a metal-coordinated ligand that is located in a suitable position to interact with the metal center. In many cases, the ligands are chosen in such a way that their geometry forces one hydrogen atom to be close to the metal atom, and thus, besides the proximity of both atoms, additional proof of the existence of a M????H bond may be required. An example of intramolecular Pt????H–N hydrogen bonding can be found in the complex [PtBr{1-C10H6(NMe2)8-C,N}{1-C10H 6(NHMe2)-8-C,H}] (1) (17). The interacting hydrogen appears in the complex as the result of the protonation of an amine nitrogen, which becomes an ammoniumtype cation and thus is quite acidic. The platinum has a formal negative charge, preserving electroneutrality and resulting in a zwitterionic complex. By using an appropriate solvent the same reaction leads to the Pt(IV) hydride [PtBrH{1C10H6(NMe2)-8-C,N}2], which is a good indication that the oxidative addition of the N–H bond may be attainable via a 3c-4e hydrogen bonding interaction. The crystal structure of 1 has been determined by X-ray 580

Table 1. Comparison of Hydrogen-Bonding and Agostic Interactions M → H–X Hydrogen-Bonding Interactions

M ← H–X Agostic Interactions

3 centers–4 electrons (3c-4e) model

3 centers–2 electrons (3c-2e) model

Electron density is donated from a filled metal orbital to the hydrogen

Electron density is donated from the X–H bond to an empty metal orbital

Favored by the electropositive character of the hydrogen; stronger with N–H and O–H, and weaker with C–H

The most usual are the M ← H–C, but M ← H–N are also known

Favored by electron-rich metal centers (Lewis bases)

Favored by electron-poor metal centers (Lewis acids)

Shift to higher frequency of the signal of the hydrogen in the 1H NMR

Shift to lower frequency of the signal of the hydrogen in the 1H NMR

No changes in the X–H coupling constant

The X–H coupling constant decreases

The typical value for the M–H–X angle is 160–170°

The typical value for the M–H–X angle is 120–130°

diffraction, and it has been found that the Pt–H distance is 2.11(5) Å and the Pt–H–N angle is 168(4)°. The proton nuclear magnetic resonance spectrum of 1 shows that the signal corresponding to the interacting hydrogen appears at a very high frequency (15.8 ppm). The presence of the signal at high frequency is a distinctive feature of any hydrogen bond and is due to the de-shielding caused by the electron density donated by the proton acceptor. This characteristic can be used to distinguish between M????H–X hydrogen bonds and M????H–X agostic interactions, given that a shift to lower frequency is observed in the case of the latter. When the proton donor is carbon, the presence of hydrogen bonding does not significantly affect the value of the JC-H coupling constant, whereas the agostic interactions cause the value of the coupling constant to decrease (15). A remarkable fact for 1 is that the signal of the interacting hydrogen is coupled with the 195Pt nucleus (nuclear spin 1/2, abundance 33.4%) and shows “platinum satellites”. The value of the coupling constant is JPt-H = 180 Hz. The existence of such a coupling is probably the best evidence of a Pt????H interaction, but not of what kind of interaction it is. Pt-H coupling has also been observed in complexes with clearly agostic interactions (15, 18). Hence, unlike the 1H NMR chemical shift, the existence of M–H coupling cannot be used to distinguish between hydrogen bond and agostic M????H interactions. Nevertheless, the value of the coupling constant can be considered as an indication of the strength of the M????H interaction regardless of its nature. Me N

Me H Pt

Me Br

N

Me H

NMe2 Pt

1

Br NMe2

2

For the complexes [PtBr{1-C 6H 4CH(R)(NR′ 2)-2C,N}{1-C6H 4CH(R)(NHR′2)-2-C,H}] (R, R′ = Me, Et; 2) (17 ), which are similar to 1, the shift toward higher frequency



in the 1H NMR is less pronounced (δ ≈ 12 ppm) and the value of JPt-H is smaller, within a range of 66–104 Hz. This is probably because the ligands are less rigid and thus there is less force on the hydrogen to locate itself above the platinum coordination plane where the 5d z orbital is located (see Fig. 4). With a neutral N–H proton donor and a cationic metal center, as in [Pt(C6H3-2,6E2)(8-acetylaminequinoline)]+ (E = PPh2, NMe2; 3) (19), the shift of the N–H hydrogen signal in the 1H NMR is smaller (δ = 13.0 ppm for E = PPh 2, δ = 10.4 ppm for E = NMe2, 9.77 ppm for the free ligand) and so are the JPt-H values (55 Hz for E = PPh 2, 33 Hz for E = NMe2). It seems that the Pt????H interaction is weaker owing to the lower electron density available in the metal and the lower electropositive character of the interacting hydrogen atom. The X-ray structure of 3 (E = PPh 2) has been determined. The Pt–H distance is 2.2(1) Å and the Pt–H–N angle is 147(9)°. While O–H is perhaps the more usual proton donor for conventional hydrogen bonding, there are very few examples of M????H–O hydrogen bonds. Complexes [NBu4][Pt(C6F5)3(hq)] (hq = 8-hydroxyquinoline or 2-methyl-8-hydroxyquinoline; 4) are the only examples studied by X-ray diffraction and NMR spectroscopy (20). The X-ray structure of 4 containing the 2-methyl-8-hydroxyquinoline ligand shows a Pt–H distance of 2.19 Å and a Pt–H–O angle of 160°. In these complexes the O–H group signal of the quinoline ligand appears in the 1 H NMR spectra at ca. 12.2 ppm—about 3.7 ppm toward higher frequency than in the case of the free ligand. Moreover, this signal shows satellites due to the coupling to the 195Pt nucleus with coupling constants of ca. 75 Hz. These values are significant but smaller than those found for complex 1 (180 Hz) or 2 (up to 104 Hz). To rationalize this difference it must be remembered that the interacting hydrogens in 1 and 2 belong to cationic fragments and are therefore expected to be more acidic, thus enhancing hydrogen bonding. On the other hand, the anionic nature of the platinum center in 4 should increase the electron density on the metal, thus favoring electron donation to the hydrogen atom.

one considers that the H is much less electropositive in the Csp –H bond than in the O–H bond. 2

R

2

H N

Me

N

C

H X

C6F5

Pt

C6F5

N

C6F5

Pt

L

X

5

6

A still smaller value of the JPt-H coupling constant (8.4 Hz) has been found for complex [PtCl2(PEt3)(benzoquinoline)], which contains the same ligand as in 5 (21). The neutral character of the metal, in contrast to its anionic nature in 5, must be responsible for these different values. In general very small 1H–195Pt couplings, if any, are found in complexes with Pt????H–C hydrogen bonds. For complexes trans-[PtX2LL′] (X = Cl, Br; L = phosphines, arsines, or olefins; L′ = pyridine-like ligands with substituents in the ortho position; 6), small 1H NMR shifts to lower field and small JPt-H (ca. 10 Hz) can be detected (22, 23). X-ray diffraction studies on one of these complexes (X = Cl, L = AsEt3, R = mesityl) show a Pt–H distance of 2.43(8) Å and a very small Pt–H–C angle of 117(4)° (23). A Ni(0) complex containing a Ni....H–N interaction has been described (24). It is [Ni(CO){NH(CH2CH2PPh2)3}][BPh4] (7), whose crystal structure shows an almost linear N–H????Ni unit (the angle is 171[6]°) and a very short Ni–H distance (1.95[9] Å). In the 1H NMR spectrum, the N–H hydrogen signal appears at high frequency (δ = 14.6 ppm).

N H P* Ni CO P* = PPh 2 7

Cl

Pt

P* P*

2-

H Me N H

Cl

Cl

Cl

N H

H

Me 8

Pt Cl Cl

O C N N

H Pt

E

Me

O

E N

H Pt

C6F5 C6F5

C6F5 3

4

As previously mentioned, C–H groups can also act as proton donors but only under certain circumstances that increase the polarity of the C–H bonds. One of these circumstances occurs when the carbon atom is in sp2 or sp hybridization. In the complex [NBu4][Pt(C6F5) 3(benzoquinoline)] (5) there is a Pt????H–C interaction (20). In the 1H NMR spectrum of 5 the signal of the interacting hydrogen is shifted toward a higher frequency (δ = 13.4 ppm; cf. 9.3 ppm in the free ligand) and shows platinum satellites with a JPt-H = 22 Hz. It should be noted that the JPt-H value is smaller in 5 than in the similar complexes of class 4, in spite of the fact that the benzoquinoline ligand constrains the position of the hydrogen on the platinum plane to a much greater extent than the 8-hydroxyquinoline type. This feature can be explained if

M????H–X Intermolecular Hydrogen Bonds To date there have not been many neutron diffraction studies on compounds containing M????H–X hydrogen bonds. A study carried out on the complex [NPr4]2[PtCl4][cisPtCl 2(NH 2Me) 2] (8) (25) shows a short intermolecular Pt????H–N interaction between one of the hydrogen atoms bonded to the nitrogen of the cis-PtCl2(NH2Me)2 fragment and the metal center of the PtCl4 fragment. The Pt–H distance is 2.262(11) Å, and the N–H–Pt angle is 164.4(7)°. A “classic” Cl????H hydrogen bond between one of the hydrogen atoms bonded to the other amine nitrogen and one chloride atom of the PtCl4 fragment (Cl–H 2.318[12] Å) is also present. Ne u t ro n d i f f r a c t i o n s t u d i e s o n t h e c o m p l e x [NHEt 3][Co(CO)4] (9) (26 ) show that the hydrogen bonded to the ammonium nitrogen points to the metal center, the Co–H distance being 2.613(2) Å and the Co–H–N angle, 180.0°. A low-temperature X-ray diffraction study on a similar compound, [(NMP)3H2][Co(CO)4]2 (NMP = N-methylpiperazine) (27) also indicates the existence of two N–H????Co interactions, with Co–H distances of 2.63(9) Å and 2.67(8) Å and N–H–Co angles of 161(8)° and 176(9)°. There are


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N–H????N interactions in the compound as well. I I CO CO CO

Co

H

H

R

Et Et

Pt

H

N

N H Et

CO

N Pt R

H

R

R

9

10

The crystal structure of complex cis-[Pt(C 6F 5) 2 (2iodoaniline)] (10) has been determined by X-ray diffraction (28). Despite the fact that the –NH2– hydrogen atoms could not be located from the density maps, their approximate position could be inferred from the geometry around the nitrogen atom. It was found that two molecules lie in parallel planes in such a way that they are linked through two Pt????H–N interactions established between one of the NH2 fragments of one molecule and the Pt atom of the other one. The Pt–H distances are 2.77 and 2.81 Å and the N–H–Pt angles are 161° and 160°. In this case it is noteworthy that despite the presence in the molecule of another potential proton acceptor, the iodine atom, interaction with the metal center is preferred. Cl

Cl C

Cl

H Ph

N

S

Pt Ph

N

S

11

An example of Pt????H–C intermolecular interaction is known. The crystal structure of the complex cis-[Pt(C6H5)2{2,2′bis(5,6-dihydro-4H-1,3-thiazine-N,N′)}] (11) (29) has shown that the hydrogen atom of a molecule of chloroform, used as crystallization solvent, is located 2.48 Å from the platinum atom, the Pt–H–C angle being 169°. Note that the polarity of the C–H bond is enhanced by the presence of the electronegative substituents bonded to the carbon atom. There is no evidence that the intermolecular interactions which are present in the solid state, as previously described, are also present in solution. Nevertheless, intermolecular M????H–ORF hydrogen bonds between [(η5–C5R5)ML2] (R = H, Me; M = Co, Rh, Ir; L = CO, C2H4, N2, PMe3) and fluoroalcohols (H–ORF) have been observed in solution by using infrared spectroscopy (30). These interactions are stronger at lower temperatures, and depending on the concentration and the acidity of the fluoroalcohol, protonation of the metal can occur. This is an important observation, since it indicates that hydrogen-bonding interactions may be the first step in oxidative additions of N–H and O–H bonds to metal centers (30), in the same way as agostic interactions have been postulated to be the first step in some activation processes of C–H bonds (15). Remember that complex 1 shows similar behavior. Concluding Remarks There are now a good number of examples of hydrogen bonds in which the proton acceptor is a transition metal. These M????H interactions are substantially different from the well-known “agostic” interactions. These kinds of interactions 582

give rise to different structural and spectroscopic features, which are compared in Table 1. M????H–X hydrogen bonding interactions are very similar to the “classic” ones. The metal has to have a filled orbital whose electrons participate, along with the hydrogen and the proton donor, in the 3-centers–4-electrons (3c-4e) system. These interactions are favored by electron-rich metal centers and thus all the M????H–X hydrogen bonds described so far involve late transition metals in low oxidation states. In the case of square planar d8 complexes, the filled metal orbital from which the electron density is donated to the hydrogen atom can be easily visualized as the dz orbital (see Fig. 4). The proton donors described so far are N–H, O–H, and C–H. In line with the electrostatic model for hydrogen bonding, the electropositive hydrogen atom and an electron pair of the metal center attract each other owing to their opposite charges. However, when the proton acceptor is platinum, theoretical studies on Pt????H–X hydrogen bonding suggest that there is a significant orbital (covalent) component in these interactions (31) and hence the existence of significant Pt–H coupling. M????H–X hydrogen bonding interaction causes a highfrequency shift in the 1H NMR signal of the interacting H. This has already been observed for conventional hydrogen bonding and it is ascribed to anisotropic de-shielding caused by the metal electron density in the proximity of the hydrogen atom and which is donated to it to establish the interaction. Nevertheless, this shift in the 1H NMR signals does not provide enough evidence of the existence of a M????H–X hydrogen bonding interaction. The added presence of coupling between the proton and the metal nucleus, when this is NMR active as is the case of 195Pt, would seem to be much more conclusive. The existence M–H coupling is not an exclusive characteristic of M????H–X hydrogen bonding interactions, since it is also found in agostic interactions. The magnitude of the coupling constant can serve as a qualitative measure of the strength of the M????H interaction. With this in mind, the results reported in the literature to date can be rationalized in the following terms. First, as in “classic” hydrogen bonds, stronger interactions are found with more polar X–H bonds. The biggest JPt–H have been measured in complexes in which the H is formally protic (1, 2), whereas with the less polar C–H bonds, very small JPt–H have been found. Second, the greater formal the negative charge that the metal center supports, the stronger the interaction is. For anionic complexes the coupling constants are greater than for neutral or cationic ones. This seems reasonable since the M→H interaction has to be favored by electron-rich metal centers. Third, the rigidity of the ligand also seems to play an important role. With rigid ligands that constrain the hydrogen atom to the appropriate position, greater couplings and therefore stronger interactions are achieved (cf. the JPt–H values for the pairs 1–2 and 5–6). Along these lines, it is noteworthy that in complexes similar to 4 and 5, [Pt(C6F5) 3L]{, where the L are pyridine- and aniline-like ligands bearing substituents in a suitable position for Pt→H interaction to be established, no coupling, and thus no evidence of interaction, was found in any case (32). The main difference between L and the quinoline-type ligands present in 4 and 5 is the rigidity of the skeleton of the ligands and the degree of freedom of movement of the hydrogen atom.


2


All these three factors usually come into play together, and the final magnitude of the strength of the M????H–X hydrogen bonding interaction is probably due to the combination of the three. In the solid state, a wide range of M–H distances, from 2 to 2.8 Å, have been found. The typical value for the M–H–X angle is in the range of 160–170°, except for some intermolecular interactions in which the geometry of the ligands causes this angle to be more acute. Nevertheless, one must be cautious in the interpretation of the X-ray and neutron diffraction data, and two situations can be differentiated. In intermolecular M????H–X hydrogen bonding interactions, such as in 8–10, there are no constraints on the position of the interacting molecules and thus hydrogen bonding occurs as an additional source of stabilization. In most cases described in the literature however, the M????H–X hydrogen bonding interaction is intramolecular, one hydrogen atom belonging to one of the ligands of the complex, the geometry of which has been chosen to force the proximity of that hydrogen to the metal center. There are some examples of complexes for which the X-ray structure has revealed relatively short M–H distances, in the range commonly found for M–H interactions, but whose 1H NMR spectra do not show M–H coupling and thus there is no conclusive evidence of the existence of the interaction in solution (32). Nevertheless, it cannot be ruled out that the weakness of the hydrogen bond in the solid state precludes its existence in solution.

7. 8. 9.

10. 11.

12. 13. 14.

Acknowledgments I thank J. Forniés, B. Menjón, and J. M. Casas for helpful discussions and the Dirección General de Enseñanza Superior (Spain) for its financial support (project PB95-0003CO2-01).

15.

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Hydrogen Bonds Involving Transition Metal Centers Acting As Proton

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