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Covalent versus Dative Bonds to Main Group Metals a Useful Distinction.

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Covalent versus Dative Bonds to Main Group Metals,
a Useful Distinction
By Arne Haaland”
Textbooks of inorganic chemistry describe the formation of adducts by coordination of an
electron donor to an electron acceptor, often using the amine-boranes, X,N + BY,, as examples. In the Lewis (electron dot) formulas of the compounds, the dative bond in H,N -+ BH,
and the covalent bond in H,C-CH, are both represented by a shared electron pair. In the
simple molecular orbital or valence bond models the wave functions of both electron pairs
would be constructed in the same manner from the appropriate sp3 type atomic orbitals on the
bonded atoms: the difference between the covalent and the dative bond becomes apparent only
after the orbital coefficients have been analyzed. This may be the reason why many structural
chemists seem reluctant to distinguish between the two types of bonds. The object of this article
is to remind the reader that the physicochemical properties of covalent and dative bonds may
be-and often are-quite different, and to show that a distinction between the two provides
a basis for understanding the structures of a wide range of main group metal compounds.
1. Introduction
In his book The Nature of the Chemical Bond, probably the
most influential book of this century on structural chemistry,
Linus P a u h g classifies chemical bonds as electrostatic, covalent or metallic. In the subsection on covalent bonds in the
introductory chapter he represents the bonding in trimethylamine oxide in the following way
In most of the adducts that we shall discuss, both donor
and acceptor are neutral molecules, but we shall include a
few species where the donor is a halide ion or the acceptor is
a molecular ion with charge 1 . Even though solvation of
ions like Mg2@or A13e,[51“secondary bonding” as encountered for instance in the solid trihalides of the heavy
Group 15 elements,161and hydrogen bonding[’] all involve
electron donor acceptor interactions, they are not within the
scope of this article.
+
2. Prototype Covalent and Dative Bonds:
Ethane and Amine-borane
and adds“] that the N-0 bond “may be considered a sort of
double bond consisting of one single covalent bond and one
ionic bond of unit strength. A bond of this type has sometimes been called a semipolar double bond. The name coordinate link has also been used, together with a special symbol
-+ to indicate the transfer of electric charge from one atom
to the other”. “We shall not find it convenient to make use
of these names or these symbols”.
Other chemists have written or edited whole books on
electron donor-acceptor adducts, e.g. E. N . Guryanova et
a1.J2I R. G. Pear.~on,[~I
and K G ~ t m a n n . [ ~ ~
In this article we first discuss the structure of simple
molecules containing covalent or dative bonds between
atoms of Groups 15, 16 or 17 and metal atoms of Groups 2
and 12- 14. Our aim is to bring together information -some
of it new, some of it quite old - which highlights the different
properties of covalent and dative bonds between the same
atom pair, for instance A1 and N or Zn and 0. In particular
we shall be concerned about bond distances. Secondly, we
wish to show how the bond distances of intermediate magnitude found in more complicated, polynuclear species may be
rationalized in terms of the magnitude of covalent and dative
bonding contributions determined by enumeration of canonical forms with localized covalent and dative bonds.
[*I
Prof. A. Haaland
Department of Chemistry, University of Oslo,
P.O. Box 1033, Blindern, 0315 Oslo 3 (Norway)
992
8 VCH firlagsgese!lschaft
mbH. 0-6940 Weinhem. 1989
H,C-CH, and H,N -+ BH, provide convenient prototypes for covalent and dative bonds: The two molecules are
isoelectronic, and, at least on paper, one species may be
converted into the other by transfer of one proton between
the nuclei of the bonded atoms.
It has sometimes been said that “the only difference between a covalent and a dative bond is where you think the
electrons come from”: When H,NBH, is formed from H,N
and BH, the two BN bonding electrons are provided by N,
but, since H,CCH, may be formed by coordination of H,C@
to @CH,, it would seem that a distinction between covalent
and dative bonds is artificial.
However, regardless of how H,CCH, and H,NBH, have
been formed, the molecules differ in the nature of the fragments obtained when the central bonds are broken:
Rupture of the dative bond in H,N + BH, yields either
two species without net charge or spin (H,N: and BH,) or
two species with both net charge and net spin (H3N*’ and
BHiO). Rupture of the covalent bond in H,C-CH,, on the
other hand, yields either species with net spin (H3C’) or
e
species with net charge (H,C: and CHT).
Minimum-energy rupture of the central bond in either molecule yields neutral species; rupture of the dative bond in
H,N -+ BH, proceeds heterolytically, rupture of the covalent bond in H,C-CH, proceeds homolytically.
Since the electron affinities of the halogen atoms are
smaller than the ionization energies of the alkali metal
0570-0833l89iO992S 02SOjO
Angew. Chem. Inf. Ed. Engl. 28 (1989) 992-1007
atoms, minimum energy rupture of the bonds in the gaseous
monomeric alkali halides-like
the covalent bond in
ethane-proceeds homolytically to yield neutral, radical species (atoms).
In the following we shall denote a bond in a neutral diamagnetic molecule as normal if minimum-energy rupture in
the gas phase or in an inert solvent proceeds homolytically to
yield neutral, radical species. The term covers both covalent
and ionic bonds as well as bonds of intermediate polarity. We
shall use this term rather than the word covalent when discussing very polar bonds and when we wish to make general
statements.
We shall define a bond in a neutral, diamagnetic molecule
as dative if minimum-energy rupture in the gas phase or in an
inert solvent proceeds heterolytically to yield neutral, diamagnetic species. The category to which a bond belongs may
then be determined experimentally or deduced from measureable quantities like ionization energies and electron
affinities.
The C-C bond rupture enthalpy of ethane at 298 K is
AH:,s = 89.8 0.5 kcal mol- '.I8] The bond rupture enthalpy of H,N + BH, may be estimated to the nearest kcalmol- from the dissociation enthalpies of seven methylated
amine boranes listed in Table 1; AH:,,(H,N + BH,) = 31.1
Table 1. Gas phase dissociation enthalpies,
(in kcal mol-'1, of the
donor-acceptor complexes of BH, and BMe, with H,N and the methylamines
Me,H, .N, n = 1-3.
\
Acceptor
BH3
BMe, [21]
31.1 k 1.0 [a]
35.0 f 0.8 [b]
36.4 k 1.0 [b]
34.8 k 0.5 [b]
13.8 i:0.3
17.6 t0.2
19.3 i: 0.2
11.6 i: 0.2
Donor
H,N
MeH,N
Me,HN
Me,N
[a] Estimated by us from the other entries in the Table. [b] Calculated from the
standard enthalpies of the reactions 2 Me,H,-,N + BH,(g) + B,H,(g)
+ 2 Me,H, .N(g) 1191 and the dissociation enthalpy of diborane 1201.
i 1.0 kcal mol-'.[*] The strength of the prototype dative
bond is thus one third the strength of the prototype covalent
bond. The difference in bond strength is reflected in the bond
distances (Table 2); the covalent C-C bond distance in
ethane is 1.533(2) 8, by GED'''] (gas electron diffraction);
the dative N + B bond distance in H,NBH, is 1.658(2) 8, by
MW[' 'I (microwave spectroscopy).[**]
While the dipole moment of ethane is, by virtue of its
symmetry, zero, the dipole moment of H,NBH,["] is 5.22 D
as compared with 1.47 D for free NH, .[81 The negative pole
is undoubtedly at the acceptor. If the electron pair was
shared equally between donor and acceptor atoms, corresponding to the transfer of one full electron, H,N@-BeH,,
the bond dipole calculated from point charges would be
8.0 D. The observed dipole moment indicates that the elec[*] This estimate is in reasonable agreement with the dissociation energy D,
obtained by SCF MO calculations with triple zeta plus polarization basis
with electron correlation Merller-Plesset calculations to the fourth order,
D, = 34.7 kcal mol-'. Correction for zero-point vibrational energies
yielded D,, = 28.7 kcal mol-' [9].
[**I
As we shall see in Section 9 a covalent N-B single bond distance may be
estimated to be 1.47 A, i.e. 0.19 8, shorter than the dative bond in
H,N + BH,.
Angew. Chem. hi.Ed. Engl. 28 (1989) 992-1007
Table 2. Comparison of the physicochemical data ofcompounds with covalent
and dative bonds.
H,C-CH,
AH:,, [kcal mol-'1
R
P
PI
M.p. ["C]
90
1.53
0
-183
H,N
+ BH,
31 [a1
1.66
5.2
124
Me$-SiMe,
74 [IS]
2.34 1161
Me,P +AIMe,
21 1171
2.53(4) [IS]
[a] See footnote to Table 1
tron pair remains more closely associated with the donor
than with the acceptor atom.
The gross atomic populations obtained by analysis of
MOs obtained with a split valence basis indicate that adduct
formation is accompanied by transfer of 0.20 eo from N to
B.["] Somewhat surprisingly, the electron density on B appears to be slightly less in the adduct than in free BH,; the
acceptor atom appears to compensate for the acquired electron density by passing a slightly larger negative charge on to
its substituent H atoms. Similarly, the donor atom N appears
to compensate for the charge transfer to B by extracting a
slightly larger negative charge from its substituent H atoms:
the charge transferred from donor to acceptor originates at
the substituents of the donor atom and ends up at the substituents of the acceptor atom. The reader may wonder
about the effect of replacing the substituent H atoms on N or
B with more or less electron-withdrawing atoms or groups.
We shall return to this point in the following section.
H,NBH, is a solid at room temperature. The melting
point['31 is about 124°C and it sublimes at 60°C/
0.001 torr.[14' The melting point of ethane is - 183 "C and
the normal boiling point is - 89 oC.[81The difference is undoubtedly due to strong intermolecular dipole-dipole attractions in the adduct.
Another pair of covalently/datively bonded molecules
that are isoelectronic and related through transfer of one
proton between the bonded nuclei is provided by Me,SiSiMe, and Me,P --+AIMe,(Table 2). In this case the strength
of the dative bond is less than a third of the strength of the
covalent bond, and the bond distance is about 10% greater.
3. Inductive Effects
Comparison of the gas phase dissociation enthalpies of
complexes of BH, and BMe, with the methylamines
Me,H,-,N, n = 1-3 (Table I), shows that replacement of
three H atoms on the acceptor atom by three Me groups
reduces the dissociation enthalpy by about 17 kcal mol- ',
i.e. by a factor of 1/2.
If the destabilization were due to increased strain following the introduction of the more voluminous Me groups,
introduction of one or more Me groups on N in
H,N + BMe, should lead to still greater destabilization. Instead, it is found that introduction of the first Me group on
N stabilizes the complex by about 4 kcal mol-' and the
second by about 2 kcal mol-', while the third Me group
leads to a small destabilization.
It might be argued that the low dissociation enthalpy of
the BMe, adducts is the result of high reorganization energy,
i.e. that it would take more energy to pyramidalize BMe,
than BH, . In fact the molecular force fields indicate that the
993
deformation energy of BH, is higher than that of BMe,.[221
We therefore interpret the strong destabilization produced
by Me groups on the acceptor atom and the weak stabilization produced by Me groups on the donor atom as an inductive effect due to the greater electron-releasing tendency of
the Me group.
If introduction of electron-releasing Me groups on the
acceptor atom destabilizes the complex, introduction of electron withdrawing CI atoms would be expected to stabilize it.
As far as we know, the dissociation enthalpy of Me,NBCI,
has not been determined. However, the gas phase dissociation enthalpies of py BMe, (py = pyridine) and
py + BCI, are available for comparison. The former,
17 kcal mol- 1,[231 is similar to that of Me,N + BMe,, while
the latter, 38 kcal mol-'[24] is similar to that of
Me,N + BH,. In borane complexes the replacement of H
atoms on B by CI atoms appears to have little effect on the
dissociation enthalpy .
It is normally assumed that bond distances may serve as
indicators of bond strength with the shorter distance associated with the stronger bond. Indeed, the N-B bond distance in Me,N + BMe,, 1.70 k 0.01 8, by MW,r251is significantly longer than the bond distances in Me,N + BH,,
I .64 0.01 8, by MW[261and 1.656(2) 8, by GED,1271and
Me,N + BCI,, 1.659(6)
and 1.652(9)
by two independent GED studies.
It is noteworthy that the bond distance in Me,N + BCI,
decreases by about 0.06 8, on going from the gaseous to the
solid state : two independent crystal structure determinations
have yielded N-B = 1.609(6) 8,f3OI and 1.575(11)w.r311
Since donor-acceptor bonds are relatively weak, they are
relatively easy to compress or elongate. The dipole moment
of Me,N
BCI, in benzene is 6.28 0.02 D.[321In the crystals, nearest neighbors are aligned with antiparallel N + B
vectors. This arrangement stabilizes greater charge transfer,
i.e. a stronger dative bond. An early X-ray crystallographic
study of our prototype adduct H,NBH, yielded an N-B
which -although very
bond distance of 1.56 0.05
inaccurate-suggests a shortening of 0.05 8, or more relative
to the gas phase.
Before proceeding to allane complexes we pause to take
note of H,BCO, which appears to be the only carbonyl complex of a main group element that is stable under the conditions of normal temperatures and pressures. Molecular orbital calculations indicate that there is considerable back
donation of BH bonding electrons into the antibonding CO
n-orbital (hyperconjugation).[' 1' The B-C bond distance is
1.534(10)
about 0.04 8, shorter than the covalent
B-CsPs bond distance in BMe,,[,'I and hence only 0.04 A
longer than estimated for a covalent BC,, bond distance.
The gas phase dissociation enthalpy, however, is only
18.8 kcal mol-',[361 or about one fifth of the mean disruption enthalpy of the covalent B-C bonds in BMe, .[151
Alane complexes appear to be as sensitive to inductive
effects as borane complexes: thus the gas phase dissociation
enthalpy of Me,N -+AICI, is AH0 = 47.5 2.0 kcal mol-'
as compared to AH:i, = 30.7 0.3 kcal mol-'
for
Me,N -+AIMe, .[371 The bond strength difference is clearly
reflected in the N +Al bond distances, 1.96(1) A by X-ray
diffraction in the chloro compound[381versus 2.10(1) A by
G E D in the methyl compound.[37]The dissociation enthalpy
--f
--f
+
+
994
of Me,N +AIH, is unknown, but the N +A1 bond distance,
2.063(8) 8, by GED,[391indicates that the compound more
closely resembles the AlMe,- than the AlC1,-complex.
In their book on the donor-acceptor bond,[*] Guryanova
and her coauthors list the dissociation enthalpies of about
700 complexes where both donor and acceptor atoms are
main group elements. The entries cover complexes where the
acceptor atom is a Group 12 element (Zn), a Group 13 element (B, Al, Ga and In), a Group14 element (Sn), a
Group 15 element (As and Sb) and a halogen atom (Br and
I). The donor atoms are Group 15 elements (mainly N and
P) and Group 16 elements (mainly 0 and S). The bond dissociation enthalpies range from 0.2 kcal mol-' to nearly
50 kcal mol- The strongest bonds to a given acceptor appear to be formed by N donors. The strongest bonds to a
given donor appear to be formed by Al acceptors, followed
by B and Ga. Starting with Group 13 the strength of the
donor-acceptor bond appears to decrease with increasing
group number of the acceptor atom. If the free acceptor
carries only hydrocarbon substituents the dissociation enthalpy is less than or equal to 30 kcalmol-' (as in
Me,N +AIMe,). If the acceptor atom carries more electronegative substituents, the dissociation enthalpy is higher,
and may approach 50 kcal mol-' (as in Me,N +AICI,).
The strength of a donor-acceptor bond between main group
elements may be as high as about half the strength of a
covalent bond, but often it is much weaker.
'.
4. Dative Bonds and the VSEPR Model
In Figure 1 we compare the structure of the free donor
Me,P[401 and the structure of the free acceptor AlMe,[411
1.9S7(3)A
I
H,C,
1.973(3) 8,
H3C
H3C'A2F3
/u
11 8.9(2)'
H3C
IP,"'CH,
?CH,
2.53(4) 8,
115.0(7)'
d<"CH3
H3C P C H ,
1.822(3) 8,
1.846(3)A
Fig. 1. Structural parameters of free Me,P [40],free AIMe, [41] and the adduct
Me,P +AIMe, [18].
with the structure of the adduct Me,P -+AIMe, .[18] In free
AIMe, the metal atom is surrounded by three bond electron
pairs and the coordination is trigonal planar in agreement
with the VSEPR (Valence Shell Electron Pair Repulsion)
model.
As the adduct is formed and a fourth electron pair is
introduced (accepted), the AI-C bonds are bent back, and
the C-AI-C angles are reduced from 120" to 113". At the
same time the bonds are slightly, but significantly, elongated
from 1.957(3) 8, to 1.973(3) 8,. Both these changes are in the
direction predicted from the VSEPR model. It should be
noted, however, that after the adduct has been formed the
C-AI-C angle remains greater than the tetrahedral angle
109.5", suggesting that the spatial requirements of the
covalent or normal bond pairs (NBP) are greater than the
steric requirement of the accepted bond pair (ABP):
NBP > ABP.
Angew. Chem. Int. Ed. Engl. 28 (1989) 992-1007
The P atom in Me,P is surrounded by three covalent bond
pairs (NBP) and one electron lone pair (LP). The structure is
pyramidal, in agreement with the VSEPR model, and the
C-P-C angle is less than tetrahedral, in agreement with the
assumption that the spatial requirement of an LP is greater
than that of a NBP.
As the adduct is formed and the lone pair is partially
removed (donated), the C-P-C angles increase from 98.6(3)"
to 103.4(8)", and there is a slight but significant decrease in
the P-C bond distances from 1.846(3) A to 1.822(3) A. Both
these changes are in the direction predicted from the VSEPR
model. After the adduct has been formed, however, the
C-P-C angle remains less than tetrahedral, indicating that
the spatial requirement of a donated electron bond pair
(DBP) remains larger than that of a covalent bond pair:
LP > DBP > NBP.
Similar changes in the structure of the donor have been
When Me,As
observed in Me3PBH,r421and Me,PBCI,
forms a complex with BH, the As-C bond distance decreases
from 1.968(3) 8,i441to 1.945 8, [451 and the C-As-C angle increases from 96.1(5)01441to 105°.1451
The structures of the adducts of AlMe, and AlCl, with
another donor, viz. Me,N are shown in Figure 2. The struc-
,
1.957(3)
, 2 064(4) A
A
groups. Formation of the adducts with AIMe, and AlCI,
leaves the C-N-C angle' unchanged or slightly decreased
while the N-C bond distance increases from 1.454(2) A in
free NMe, to 1.473(3) 8, in the complex with AlMe, and to
1.516(12) A in the still stronger adduct with AICI,.
It appears that while the structural changes in Me,P could
be rationalized in terms of VSEPR, the structural changes in
Me,N on adduct formation are in the opposite direction and
must be rationalized in terms of steric repulsion between the
atoms and groups bonded to N.[*]
Consideration of steric repulsion between the ligands bonded to the acceptor atom (Al) leads to the prediction that
A1-Y bonds should be elongated and Y-AI-Y angles reduced
when the adduct is formed. On the acceptor atom VSEPR
and steric interactions reinforce one another, on the donor
atom they are opposed. When the donor atom is small (N or
0),steric factors may prevail.
According to Gutmann's rules on bond length variati~n,'~'
adduct formation should always lead to elongation of bonds
to both donor and acceptor atoms. This rule appears to be
generally valid for bonds to the acceptor atoms, but is broken for bonds to the donor atom when Me3P and Me,As
form adducts with boranes or allanes. We presume that it
may be broken in other cases when the donor atom is large
enough to make steric interaction between the substituents
unimportant.
5. 2: 1 Arnine-Alane Complexes
2.121(4) 8,
.Aid c1
C!
c1
102.3(3)"
2.099(10) 8,
Both Me,N +AlH3 and Me,HN +AlCl, will add a second mole of the electron donor. The resulting coordination
around A1 is trigonal-bipyramidal with the donor atoms in
axial positions as expected from the VSEPR model with the
additional postulate that NBP > ABP.
The N +A1 distances in the 2: 1 adducts (Fig. 3) are considerably longer than in the 1:l adducts, indicating satur-
108.3(2)"
d<"CH3
H3C k H 3
1.454(2)8,
Fig. 2. Structural parameters of free AIMe, 1411, free AICI, [46], free Me,N [47]
and the adducts Me,N +AIMe, [37] and Me,N 4AlC1, [48] (the AI-N bond
length was determined X-ray crystallographically 1381.
,N,,,
CH,
CH3
\
ture of the acceptor (AIMe,) in the adduct with Me,N is very
similar to the structure in the adduct with Me,P and may be
rationalized in the same manner. Monomeric AlCl, is trigonal planar in the gas phase with Al-Cl = 2.064(4) A.r461On
formation of the adduct the CI-AI-Cl angle is reduced from
120" to 114" and the AI-CI bond distance increases to
2.121(4) A; again the changes are those predicted from the
VSEPR model.
We now turn our attention to the structure of the donor:
In free Me,N the N atom is surrounded by three CBP and
one LP. According to the VSEPR model the molecule should
be pyramidal and the C-N-C angle less than tetrahedral. The
C-N-C angle is, however, found to be slightly greater than
tetrahedral, 110.6(6)' by GEDr4'l and 110.9(6)O by MW,[491
presumably because of steric repulsion between the Me
Angew. Chrm. I n ! . Ed. Engl. 28 (1989) 992-1007
H3C"L
\ CH3
CH3
Fig. 3. The structural parameters of (Me,HN),AICI, [S3] and (Me,N),AIH,
determined by X-ray diffraction [54] and gas-phase electron diffraction [ S S ]
ation of the acceptor: in (Me,HN), AlCl,,[53' N +Al =
2.059(3) A compared to 1.96(1) A in Me,NAICl,[38J
(Fig. 2). In (Me,N),AlH,,r547551N +Al = 2.18(1) A compared to 2.063(8) A in the 1 :1 adduct. The N +AI bonds in
[*] This difference between N and P appears to be carried over to tetraalkylarnmonium and -phosphonium ions: The mean N-C bond distance in
[Me,N]F. 4 H,O is 1.499(2) A [50], 0.04,A longer than in free (CH,),N,
while the mean P-C bond distance in [Me,P][Cu,CI,] is 1.779(3) A [Sl],
0.06, A shorler than in Me,P. The average of 319 N-C* bond distances
(C* = tetrahedral C atom carrying only C and H as additional substituents)
in NC:-type ammonium ions is 1.510 A, 0.04, A longer than the average of
1042 N-C bond distances in NC:-type amines, while the average of 35 P-C*
bond distances in C,PC*-type phosphonium ions is 1.800 A, 0.04, A shorter than the average of 23 bond distances in C,PC* phosphanes 152).
995
(Me,HN),AlCI, are about 0.12 8, shorter than in
(Me,N),AIH,, again indicating that AICI, is a better acceptor than AlH,. The N -+A1 bond distances in the 1 : 1 adducts
indicate that AlMe, is a poorer acceptor than AIH,; indeed
Me,NAlMe, will not add another mole of the base: it appears that 2: 1 adducts are as sensitive to inductive effects as
the 1 : 1 adducts.
Both VSEPR and steric considerations suggest that covalent AI-Y bonds should be further elongated on addition
of the second donor. Indeed, in (Me,HN),AlCI,,
A1-Cl = 2.180(1)
as compared to 2.121(4) i%in
Me,NAICl,
We are not aware of any observation indicating the existence of 2: 1 amine-borane adducts.
6. Some Adducts of Beryllium, Magnesium, Zinc,
and Mercury
Quinuclidine forms 2: 1 adducts with both dimethylberyllium and dimethylmagnesium. The coordination geometries
around the metal atoms as determined by X-ray crystallog r a p h ~ [571
~ are
~ , shown in Figure 4.[*] The main features are
1.83 8,
2.19
A
I .98 8,
231A
both X-ray and GED.[631The Zn-C bonds are about 0.05 8,
longer than in d i p r o p y l z i n ~ , [the
~ ~dative
]
N + Zn bonds are
0.50 8, longer than the covalent Zn-N bonds in Zn[(NSiMe,),], .I6'' The weakness of the donor-acceptor interaction
is indicated by the large C-Zn-C valence angle and the large
N + Zn r.m.s. vibrational amplitude found in the GED
study, 1 = 0.17(3) A, two or three times greater than the
amplitude of the Zn-C bond. The donor-acceptor bonds
appear to be about 0.08 8, shorter in the solid phase than in
the gas phase.
The structure of the 2:l complex of pyridine with
ZnC1,[651is shown in Figure 5. The N + Zn bond distance is
about 0.25 8, longer than the covalent bond distance in
Zn[N(SiMe,),], but 0.25 8, shorter than the N + Zn bond in
2; the latter difference is presumably an inductive effect. The
Zn-Cl bond is 0.1 5 8, longer than in gaseous monomeric
ZnC1, .Iti6]
In Figure 5 we also show the coordination sphere of the
mercury atom in the related adduct of bis(2-pyridyl) disulfi2.22
A
2 34
A
Fig. 4. Structural parameters of the 2: 1 adducts ofquinuclidine with Be(CH,),
and Mg(CH,), [56, 571.
in agreement with the principles already outlined: BeMe, is
linear in the gas phase with R(Be-C) = 1.70
coordination of two donor atoms decreases the C-Be-C angle to
1 18"and increases Be-C to 1.83 8,. In monomeric, base-free
bisneopentylmagnesium the C-Mg-C angle is 180" and
R(Mg-C) = 2.13 8,;["lin theMgMe,adduct thevaluesare
129" and 2.19 8,.
In each complex the metal atom is surrounded by two
covalent electron bond pairs and two accepted bond pairs;
the structure is distorted tetrahedral. Even though the datively bonded quinuclidine ligands are more bulky than the
covalently bonded Me groups, the C-M-C angles are considerably larger than the N-M-N angles, in agreement with the
postulate that NBP > ABP.
The interaction enthalpy between 2,2'-bipyridine and dimethylzinc in toluene is 17.6 kcal mol-'.1601 Since 2,2'-bipyridine presumably acts as a bidentate ligand, this corresponds to a mean N + Z n dissociation enthalpy of
9 kcal mol-'. The Zn-N bonds in the amides Zn[N(SiMe,),], are assumed to be single bonds with negligible Kcomponents.[611The bond energy term[621is 49 kcal mol- I ,
i.e. the strength of the covalent bond appears to be four or
five times that of the dative bond.
We are not aware of the structure of any simple adduct
between a dialkylzinc and an N-centered donor, but the
structure of Zn[(CH,),N(CH,),], 2 has been determined by
[*] In the following we omit estimated standard deviations of less than 0.01 8,
or I".
996
Fig. 5. Structural parameters of ZnCI, 2 pyridine [65]and HgCI, . bis(2-pyridyl) disulfide 1671 in the crystal.
de with HgC1, in the solid phase.[671Mercury compounds are
known as poorer acceptors than their zinc analogues. The
weakness of the interaction is indicated both by the large
N + Hg bond distance, more than 0.50 8, longer than the
covalent Hg-N bond in the amide Hg[N(SiMe,),], ,[681 and
by the large Cl-Hg-CI angle. The valence angles at the Hg
atom are such that a description in terms of sp-hybridization
with electron donation into pure p-orbitals may be more
appropriate than a discussion in terms of sp3-hybridization.
The elongation of the Hg-C1 bond compared to that in gaseous monomeric HgCI, is 0.09
The most interesting N --t Zn adduct may be bis(1,Cdihydropyridin-I -yl)-bis(pyridine)zinc, 3, which is formed from
3
ZnH, and four moles of pyridine.[". 711 In this complex we
see for the first time the difference between covalent and
dative bonds between the same atom pair demonstrated in
the same molecular unit.
The covalent Zn-N bonds are about 0.1 5 8, longer than in
Zn~(SiMe,),], : this elongation is similar to the elongation
Angen,. Chenz. Int. Ed. Engl. 28 (1989) 992-1007
of the bonds in ZnC1, on coordination of two pyridine molecules. The covalent Zn-N bonds in 3 are, however, still
0.16 8, shorter than the dative bonds. The bulks of the four
ligands are very nearly equal, and, as expected from VSEPR
considerations, the angle subtended by the covalently bonded N atoms (124") is considerably larger than that subtended by the datively bonded N atoms (98").
is then the sum of the N + Si bond dissociation enthalpy and the barrier to rotation of the SiF, group in the uncoordinated intermediate.
Comparison of the axial N -+ Si bond distances found in
the crystal structures of the related compounds 5[771and
6[781indicates that the dative N + Si bonds are even more
F
I
7. Some Adducts of Silicon
The 1:l adducts of SiF, with the amines Me,H,-,N,
n = 0,1,2,3, in argon matrices have been investigated by in731 The coordination around Si was
frared spectrosc~py.'~~.
found to be trigonal bipyramidal with the amine occupying
an axial position. A structural chemist has
"This conclusion is remarkable. Of the huge range of fivecoordinate derivatives known of P and s, all examples with
only monodentate groups have the axial sites occupied by
the most electronegative ligand . . .". In other words, why
does the N atom in Me,NSiF, occupy an axial position when
the N atom in Me,NPF,'751 is equatorial (Scheme l)?
F
F
F
sensitive to inductive effects than N +A1 or N + B bonds:
replacement of three F atoms on Si by three Me groups
increases the dative N + Si bond distance by 0.73 A!
The covalent Si-N bond distance appears to increase by
0.02 8, on going from 5 to 6,but the combined uncertainty
is so large that it is doubtful whether the difference is real.
A more accurate estimate of the effect on covalent bond
distances when three F are replaced by three Me groups, is
obtained by comparing the Si-F bonds in SiF,[791 and
Me,SiF["] and the Si-C bonds in MeSiF,'"I and Me,Si:[8'1
in both cases the bond distances increase by 0.05 A. The
difference between the N -+ Si bonds in 5 and 6 is thus more
than an order of magnitude greater than the elongation expected for covalent bonds.
Scheme 1. Structures of Me,NSiF, and Me,NPF,
8. Influence of the Donor Atom and Net Charge
We suggest that the contradiction is removed if a distinction is made between covalently bonded ligands (F and
NMe,) and datively bonded ligands (NMe,). If the distinction is made, VSEPR considerations combined with the
postulate NBP > ABP lead to the observed structures for
both molecules.
The crystal structure of 2-{[methyl(trifluorosilyl)amino]methy1)pyridine 4 shows that the coordination of Si is trigo-
4
nal bipyramidal.[761Two F atoms and the covalently bonded
amide N atom occupy the equatorial sites, while the axial
sites are occupied by the third F and the datively bonded
pyridine N atom. The Si-F,, and Si-Fa, distances are very
similar, 1.603(4) 8,and 1.621(5) A, respectively. The equatorial Si-N distance is 1.70 A, a normal value for a covalent
Si-N bond, the dative, axial N + Si distance is 1.97 8,.
I9F-NMR spectra show that axial and equatorial F atoms
undergo rapid exchange at room temperature.[761There is
evidence that the exchange proceeds via dissociation of the
dative N + Si bond followed by rotation of the SiF, group
and recoordination of the pyridine N atom. The activation
energy (13.2(9) kcal mol-') determined by line shape analyAngew. Chem. Ini. Ed. Engl. 28 (1989) 992- 1007
8.1. A Comparison of Donor Atoms
Trimethylaluminum forms stable complexes with a wide
range of electron donors. Some D +A1 bond distances and
gas phase dissociation enthalpies [Eq. (a)] are listed in
Table 3.
Table 3. AI-D bond lengths and gas phase dissociation enthalpies of some
complexes of AIMe, and electron donors D.
Donor D
R [A1
AH:,%[kcal mol-'1
Me,N [37]
Me,P [17, 181
Me,O [17, 821
Me,S [83, 841
Me,Se [SS]
2.10
2.53(4)
2.01(2)
2.55(2)
31
23
22
18
16
-
Complexes of trimethylaluminum with alkyl fluorides or
chlorides have not been isolated; the donor-acceptor interaction is obviously not strong enough to compensate for the
enthalpy required to produce one mole AlMe, by dissociation of the dimer, viz. 10 kcal m ~ l - ' . [ It
~ ~should
]
be noted,
however, that Et,SiF forms a stable complex with AIEt, .[*'l
The enthalpy required to break the dative F +A1 bond is
997
obviously greater than the enthalpy gained by dimerization
of one mole of Et,AI (9 kcal mol-'[*']).
BH, will only form complexes with electron donors in
those cases when the donor-acceptor interactions are strong
enough to compensate the diborane dissociation enthalpy.
No such limitations exist for BF, which is monomeric. Gas
phase dissociation enthalpies are available for the complexes
with Me,N, Me,O, Me,P and Me$ and the structures of the
former two have been determined by GED (Table 4).
Table 4. B-D bond lengths and gas phase dissociation enthalpies of some complexes of BF, and electron donors.
[A1
Donor D
R
Me,N [88, 891
Me,P 1921
Me,O 190, 911
Me$ [93]
1.61
~
1.75(2)
-
27
19
13
4
the sublimation and dissociation enthalpies. The experimental values are[941 45 kcal mol-' for D = Me,P and
20 kcal mol- ' for D = Me,As. Subtraction of the dissociation enthalpy of Me,PBF, from the value for Me,P yields
a sublimation enthalpy of 26 kcal mol-' for the adduct.
Since the sublimation enthalpy of Me,AsBF, is expected
to be greater than that of the P-analogue, it follows from
the value AH: = 20 kcal mol- ' for Me,As that the dissociation
enthalpy
of
the As-adduct
is negative
(less than - 6 kcal mol-'). It seems that in this case intermolecular forces in the crystal are sufficient to stabilize
an adduct where the donor-acceptor interaction is repulsive
in the gas phase!
The structures of the adducts HC1BF,1951and ArBF,1961
have been determined by microwave spectroscopy
(Scheme 2).
p = 0.18 D
Scheme 2. Structural parameters of HCIBF, and ArBF,
These species are commonly regarded as van der Waals
complexes, but may also be described as weak donor-acceptor adducts. Comparison of the Ar + B and C1 -+ B bond
distances, and comparison of the dipole moment of ArBF,
with the dipole moment component pa. along the C1 + B
bond in HCIBF,, indicates that Ar is the weaker donor. HCI
and Ar are isoelectronic; on paper, HCI may be formed from
Ar by withdrawing one proton from the nucleus. The reader
may wonder what effect it would have on the donor-accep998
8.2. Cationic Complexes
A H & [kcal mol-'1
The enthalpy A H : of reaction (b) is equal to the sum of
pa = 0.48 D
tor interaction if the proton was completely removed to leave
[CIBF,]@.We shall return to such species in Section 8.3.
We have already pointed out that the strongest dative
bonds to a given main group acceptor atom are formed by N
donors. Beginning with Group 15 the strength of the donoracceptor bond appears to decrease with increasing group
number of the donor atom. With B and Al acceptors the
strength of the interaction decreases monotonically with increasing principal quantum number of the valence electrons
of the donor atom, but this does not appear to be true for all
acceptors.
Since the introduction of electron-withdrawing substituents on the acceptor atom often leads to a significant
increase in the dissociation enthalpy of a donor-acceptor
adduct, it is hardly surprising that cationic acceptors form
stronger dative bonds than related neutral molecules.
The tetramethyl derivatives of Si, Ge or Sn do not appear
to interact with electron donors. Introduction of a CI atom
on Sn leads to the formation of a stable 1 : l adduct
with pyridine, the dissociation enthalpy in benzene being
11.0 f 1.0 kcal m ~ l - ' . ~ ~ The
']
interaction enthalpies of
Me,SiCI or Me,GeCI with pyridine in the same solvent are,
however, of the order of - 2 kcal mol-' or I ~ S S . ~ ~ ' ~
In striking contrast the dissociation enthalpies of the
adducts of the cations Me,Me with the donor H,O determined by high pressure mass spectrometry are 30.1 f
1.9 kcal mol-' for M = Si,[981 28.6 0.5 kcal mol-' for
M = Ge,[991and 25.7 f 1.0 kcal mol-' for M = Sn.['OO1[*l
When Me,M@ is formed from Me,M by removal of CHF ,
a coordination site becomes vacant both in a steric sense
since the coordination number has been reduced from four
to three, and in an electronic sense since the number of electrons in the valence shell of M has been reduced from eight
to six. At the same time the formal charge on the metal atom
has increased from zero to + 1.
Me,AI, Me,Ga and Me,In are isoelectronic and isoleptic
with Me,M@, M = Si, Ge and Sn respectively. Comparison
of the dissociation enthalpies of related adducts (Table 5)
shows that the introduction of a unit positive charge on the
acceptor atom doubles or triples the dissociation enthalpy.
Comparison
of the dissociation
enthalpies of
Et,O -+ GaMe,, Et,O -+ GaCI, and [H(Me)O -+ GeMe,le
indicates that in this series the stabilization accompanying
the introduction of a net positive charge on the acceptor is
twice as large as the stabilization accompanying the introduction of three electron-withdrawing CI atoms.
The dissociation enthalpies in Table 5 also show that replacement of an H on the donor atom with a more electronreleasing methyl group invariably leads to stabilization of
the adduct, and that the first methyl group introduced has
the largest stabilizing effect.
Before we go any further, we note that while the acceptor
ability of the cations Me,M@ decreases in the order
[*I
In the absence of evidence to the contrary we assume that low-energy rupture of M-0 bonds in [H,OM Me,]@ would proceed heterolytically to yield
H,O and Me,M@,i.e. that the 0-M bond is dative according to the definition given above.
Angew. Chem. Int. Ed. Engl. 28 (1989) 992-1007
Table 5. Gas phase dissociation enthalpies, AH:,$ of cationic and neutral
donor-acceptor adducts.
~-
Donor
Acceptor
AH:,, [kcal mol-'1
HZO
MeOH
EtOH
Et,O
Et,O
H,O
MeOH
Et,O
Et,O
HzO
MeOH
NH,
MeNH,
Me,NH
Me,N
Me,N
Me3SIe [98]
Me,Si@ [98]
Me,Si' [98]
Me,Sie [98]
Me,AI [a]
Me3Gee [99]
Me3Gee [99]
Me3Ga [loll
CI,Ga [lo21
Me&@ [loo]
Me,Sne IlOO]
Me,%@ [loo]
Me,Sne [IOO]
Me,Sne [loo]
Me,%@ [loo]
Me& [I031
30.1 f 1.9
39.2
42.0
44.2
20.2 f 0.2
28.6 +_ 0.5
35.0
12.1 f 0.2
26.6
25.7
32.6 f 1.0
36.9
42.1
44.2
45.6
19.9 f 0.5
In Figure 6 we compare the average bond distances in the
The two Ga-C1
ions [Me4-nGaC1,] (n = 1 -4).1107,111,1121
bond distances in [Me,GaCl,]@ are indistinguishable, the
normal : dative 0 : 1
1:l
normal : dative 2 : 1
3:l
[a] in hexane 1171
Sie > Gee > Sn@,the normal order for neutral molecules is
considered to be Sn > Ge > Si.1971
Both Me3Sie and Me3Gee may be prepared in solut i ~ n . f ' ' ~I' . The structure of 7,the adduct of Me3Sie with
pyridine, has been determined by X-ray diffraction.['061The
T
Fig. 6. Average Ga-C and Ga-Cl bond distances in the ions
[(CH,), .GaCIJe (n = 1, 2, 3 or 4) as determined by X-ray diffraction with
[Me,AsIe or [Me,Sb]@ as counterion [107, 111, 1121 and the ratio of covalent
to dative contributions to the Ga-CI bonds.
mean value (2.28 A) is 0.10 A shorter than the Ga + C1'
bond distance in [Me,GaCl]@ and 0.18 8, longer than the
terminal Ga-Cl distance in Ga,Cl,.[' 'I The bonding may be
described in terms of the canonical forms A and B, where the
1.86A
0
mean Si-C distance is not significantly different from the
mean Si-C distance in the neutral five-coordinate complex 6,
but the dative N + Si distance is more than 3/4 A shorter:
again we find that the dative bond is more susceptible to
inductive/charge effects than covalent bonds. It may be noted that even the very short dative bond in 7 is 0.1 1 A longer
than the covalent, equatorial Si-N bond in 6.
8.3. Anionic Complexes
We now turn our attention to adducts where the donor is
a negatively charged, monatomic ion. The structure of the
anion of [Me,As][Me,GaCl] is distorted tetrahedral with a
mean Cl-Ga-C angle of 103" and a mean C-Ga-C angle of
1 15".f1071
The mean Ga-C bond distance (2.05(2) A) is not
accurately determined, but appears to be 0.08(2) A longer
than in free Me,Gar'oS1 and 0.05(2) A longer than in the
adduct Me,NGaMe,.['09' The Ga-Cl distance (2.38 A) is
0.28 A longer than the terminal bond distance in Ga,Cl,
(2.10 A),'' 'I which may serve as a reference value for a
single normal Ga-CI bond. If it is assumed that the electron
affinity of C1 is greater than that of Me3Ga, it follows that
minimum-energy rupture of the Ga-Cl bond in [Me,GaC1Ie
would proceed heterolytically to yield Me,Ga and Cl0, i.e.
that the Ga-C1 bond is dative, Ga + Cle, according to our
definition. In the following we adopt this description which
is in good agreement with the observed bond distances and
angles.
Angen. Chem. I n [ . Ed. Engl. 28 (1989) 992-1007
CH3
I
H3C-Ga-CI
t
CH3
I
H,C-Gat CI'
I
CIR
Cl
A
B
Ga atom forms three normal and one dative bond in each.
Delocalization (resonance) would yield equal contributions
of dative Ga + CIo and normal Ga-C1 bonding to each of
the two Ga-Cl bonds in the real ion. We note that the observed Ga-Cl distances are reasonably close to the mean of
dative Ga t Cle and normal Ga-Cl bond distances
[(2.10 8, + 2.38 A)/2 = 2.24 A] and choose, therefore, to describe the Ga-Cl bonds as intermediate between pure normal
and pure dative bonds. This description is consistent with the
expected fragmentation behavior of the ion; one Ga-CI
bond (the first) is expected to yield Me,GaCl and Cl0, the
other (the second) to yield Me,Ga and C1.
Similarly, the bonding in [MeGaCI,]@ may be discussed
in terms of three canonical forms, each with three normal
Ga-C or Ga-Cl bonds and one dative bond (Ga + Cle).
Delocalization yields normal to dative bonding contributions equal to 2: 1 , and the Ga-Cl bond distance in the real
ion may be estimated from (2 x 2.10 A + 3.28 A)/
3 = 2.19 A, versus an observed value of 2.22 A.
Finally, the Ga-CI bonds in [GaC141e may be described in
terms of 3:l contributions of normal and dative bonding,
and the bond distance estimated from (3 x 2.10 8, + 2.38 A)/
4 = 2.17 A, equal to the observed value.
All things considered, we feel that the agreement between
estimated and observed Ga-CI bond distances is as good
as can be expected for a linear, two-parameter model. If
999
the standard bond distances R(Ga-Cl) = 2.12 8, and
R(Ga t Cle) = 2.40 A had been chosen, the four bond distances in Figure 6 and the terminal bond distance in Ga,CI,
would all have been predicted with a deviation of 0.02 A or
less.
The crystal structure of [Ph,PNPPh,][CIBH,] has recently been determined by both X-ray and neutron diffraction
(Scheme 3).[1'31The [CIBH31e ion is pyramidal with a mean
B-H bond distance of 1.20 8, and a CI-B-H angle of 105".
"It has previously been suggested that aluminum-nitrogen
bonds that fall at about 1.78 A are in effect double bonds,
with dative nitrogen lone pair to aluminum 3d x-donation,. . .". We note that the Al-N bond distance calculated
from the modified Schomaker-Stevenson (MSS) rule
[Eq. (c)]~'"~is 1.77 A, and suggest that the shorter A1-N
bond should be described as a single covalent bond.
R(A-B)
= rA
+ rB - 0.085 IzA - ~
~
1
~
.
~
(4
The trisamide AI[N(SiMe,),], is monomeric with trigonal
planar configuration at A1 and R,,-, = 1.78(2) A.r1I91 In
this species, where A1 is tricoordinate, the AI-N bonds may
2.00 A
be shortened by dative x-bonding. Comparison with the
bond distance quoted above indicates that the effect is modCI
est.
Scheme 3. Structural parameters of [CIBH,]@
Aluminum trisisopropoxide Al(OiPr), is tetrameric in the
solid phase.['201The central A1 atom is connected to three
terminal A1 atoms through double alkoxide bridges. Each
The authors remark: "The observed bonding parameters are
terminal A1 atom is surrounded by two bridging and two
quite normal, except that the B-Cl distance on the anion is
terminal isopropoxide groups. The terminal A1-0 distances
anomalously long", and add that this "can be ascribed to
(1.71 A) are in good agreement with the covalent bond disincorporation of approximately 19 % (B,H,)@ into the
tance calculated according to the MSS rule[1181(1.72 A).
( C I B H , ) ~site".
This
covalent bond distance is 0.30 8,shorter than the dative
We would suggest that the observed bond distance of
0
+A1
bond distance in Me,O +AIMe, (2.01(2) A).
2.00 8,may be quite appropriate for a dative bond from C1°
The
structure
of the monomeric aminoborane, H,NBH,,
to B: the difference of 0.38 A between the dative Ga-CI bond
generated
by
pyrolysis
of diborane and ammonia at 500 "C,
in [Me,GaCl]@ and the B-CI bond in [CIBH,]@corresponds
MW spectroscopy.[I2'] The molehas
been
determined
by
to the difference in the bonding radii of Ga and B. SCF MO
cule has a planar ethene-like structure with a N-B bond
calculations on [CIBH,le with a 6-31G* basis set yields an
distance of 1.391(2) 8,. This bond may be described as havoptimum structure close to the experimental, and a binding
energy relative to BH, and Cle of 19.9 kcal r n ~ l - ' . ~In' ~ ~ ~ ing partial double bond character; a covalent o-bond plus a
dative x-bond. What is the effect of the x-bonding on the
the complex HC1- BF, (see Section 8.1) the B-Cl bond disbond distance? The author of a recent review['221surveyed
tance was 3.17 8,; as expected CIe is a better electron donor
the
(dative) bond distances in amine-boranes and adopted
than HC1.
N-B
= 1.58 8, as reference value for a single (covalent) bond
The covalent B-CI bond distance in BCl, is 1.74
distance.
This value would indicate that dative x-bonding in
The [BC1,Ie ion in [NS,Cl,][BCl,] is slightly distorted tetraH,NBH,
shortens the BN distance by 0.19 A, which is simhedral with all B-CI bond distances close to the mean value
ilar
to
the
difference between single and double CC bond
of 1.84
Since B, like Ga, can form three covalent
distances.
bonds to C1 and one dative bond to CI0, delocalization
If the dative bond distance in an amine-borane is to serve
yields covalent to dative bonding contributions equal to 3 : 1.
as
expectation value for a single covalent bond distance in
The estimated bond distance [(3 x 1.74 8, + 2.00 A)/4 =
H,NBH,, the logical choice would be the bond distance in
1.83 A] is in good agreement with the observed value.
H,NBH,, B-N = 1.66 8,.r111 Adoption of this value indicates that dative x-bonding in H,NBH, shortens the BN
bond by 0.27 A, which is considerably more than the differ9. Covalent AI-N, Al-0 and B-N Bond Distances
ence between single and double CC bond distances.
The bonding in H,NBH, has been studied by ab initio
'
I
molecular
orbital calculations with a double zeta (DZ) basis
The two A1-N bond distances in 8 differ by 0.19
(Scheme 4).[1231Complete structure optimization yielded an
The longest bond, from A1 to the four-coordinate N atom,
has the same length as the A-N bond in Me,N +AICI, and
1.77 %, H,C'CH3
I
CI,, \
CI -2 A1",
BE = 33 kcal mol-'
,,,
/
1.35 A
1,
\
',,
,c-c
l.SOA/
\
A E = 6 5 kcal mol-'
must be described as a dative bond. About the shorter bond,
from A1 to the three-coordinate N atom, the authors write:
1000
Scheme 4. Calculated structural parameters for planar and orthogonal
H,NBH, and C,H,.
Angew. Chem. Inr. Ed. Engl. 28 (19x9) 992-1007
equilibrium B-N bond distance of 1.378 8,, in good agreement with the experimental value. When the n-bond is broken by twisting the BH, fragment through 90” into an orthogonal form, the B-N bond distance increases to 1.469 8,
and the configuration at N becomes pyramidal.
The energy of this relaxed orthogonal form is
33.3 kcdl mol-’ above the energy of the planar equilibrium
structure. This energy difference appears reasonable when
compared to the barrier to internal rotation in Me,NBMePh
obtained by line shape analysis of ‘H NMR spectra,
26.3 kcal mol - .[I 241
MCSCF calculations on the planar and orthogonal forms
of ethene[’251 (Scheme4) yield a a-bond energy of
65 kcal mol-’, in good agreement with experimental estimates. It appears that the dative n-bond energy in H,NBH,
is about half the covalent n-bond energy in ethene. The C-C
distance in the orthogonal form of ethene is that expected for
a single bond between two sp2 hybridized carbon atoms.
In the trisborylamine 9 one B,N-BS, fragment adopts an
orthogonal conformation, presumably for steric rea-
sons.[’ 261 The corresponding B-N bond distance is l .47 A, in
agreement with the theoretical value. In the following we
shall adopt this as a standard value for a single covalent B-N
bond distance. If this value is accepted, it follows that nbonding shortens the B-N bond in H,NBH, by about
0.07 A, a third of the difference between single and double
CC bond distances.
Coordination of an increasing number of n-donors (amido
groups) to the same n-acceptor (B), would be expected to
lead to saturation effects. Indeed, the BN bond distance in
B(NHMe), is 1.432(2)
1281
Hexagonal boron nitride (a-BN) has a graphite-like structure consisting of parallel, planar BN sheets.[1291Each B and
N are trigonal planar, the B-N bond distance is 1.446 A, and
the perpendicular distance between the layers is 3.30 as
compared to 3.35 8, in graphite. The B-N bond distance is
slightly, but significantly, longer than in B(NHMe), , perhaps because each N atom is coordinated to three rc-acceptors rather than one.
According to Gmelins Handbook of inorganic Chemist r y : [ 1 2 9“Die
1
Natur der B-N-Bindung ist umstritten. Einerseits Ihnelt die BN-Struktur weitgehend der Graphitstruktur, die Bindungslangen C-C und B-N sind nur wenig
voneinander verschieden, und die zwischenatomaren Krafte
beider Substanzen liegen in der gleichen GroBenordnung.”
“Die BN-Bindungslange in der Schichtebene betragt 1.446 A
und ist damit sehr vie1 kiirzer als die Summe der kovalenten
Radien bei Einfachbindung (1.57 8,)”. “Es liegt also nahe,
fur BN eine Resonanzstruktur wie in Graphit anzunehmen.
Andererseits weichen gewisse Eigenschaften des BN stark
von denen des Graphits ab. So hat BN eine geringe elektrische Leitfahigkeit, ist farblos, und seine diamagnetische
Suszeptibilitat ist gering, wahrend die des Graphits abnorm
groB ist. Diese Eigenschaften des BN lassen sich durch das
Vorhandensein von Einfachbindungen erklaren.” [*I
We note that the B-N bond distance in a-BN is in fact
close to our estimate of a covalent single bond distance,
1.45 8, vs 1.47 8,; the contradiction has been removed.
A rough estimate of the strength of a single covalent B-N
bond may be obtained by calculating the energy of atomization of hexagonal boron nitride (at 25 “C) [Eq. (d)][1301
AEf,(BN,hex)
=
=
+
AHP(B,g) AHF(N,g) - AHP(BN,hex) - 2 RT
307.3 kcal mol-I.
(4
and assuming the energy of atomization to be equal to the
energy required to break three o-and one n-bond [Eq. (e)].
AE:,(BN,hex)
=
3 E(B-N)
+ E,(B-N)
(e)
E,(B-N) = 33 kcal mol-’ yields the mean B-N single bond
energy of E(B-N) = 91 kcal mol-’. An independent,
though equally rough estimate based on the energy of atomization of cubic boron nitride (p-BN) yields E(B-N) =
92 kcal mol-’.[**]
Both these values are overestimates; “delocalization energies” or “resonance energies” have been disregarded. We
believe E(B-N) = 88 i 2 kcal mol-’ to be a more realistic
estimate. In any case there is little doubt that covalent B-N
bonds are as strong as or stronger than covalent C-C bonds:
the mean C-C bond energy derived from the energy of atomization of diamond is 85 kcal mol- The characteristics of
covalent and dative B-N bonds are listed in Table 6.
’.
Table 6. Comparison of covalent and dative B-N bonds
Covalent [a]
~~~~
E [kcal mol
~
R [%.I
I
]
Dative [b]
~
88 2
1.47 i 1
31
1.66
[a] Derived from comparison data. [b] Determined on H,N
+ BH,
A.[1273
A n g e w Chem. I n f . Ed. Engl. 28 (1989) 992-1007
10. Some Polynuclear AlN Compounds
The trisamide Al(NMe,), is dimeric in the solid
phase.[1321Each Al atom is surrounded by two bridging and
two terminal NMe, groups. The terminal AI-N bonds have
a length of 1.81 8, and may be described as single covalent
bonds. The bridging AI-N bonds are 1.97 8, in length and
[“I
[““I
“The nature of the B-N bond is controversial. On the one hand the BN and
graphite structures are very similar, the difference between the C-C and
B-N bond distances is very small, and the interatomic forces in the two
substances are of the same order of magnitude”. “The B-N bond distance
in the sheets is 1.446 8, and thus much shorter than the sum of thecovalent
radii 1.57 8,”. “It is therefore reasonable to assume a resonating (single
and double bond) structure as in graphite. On the other hand, some properties of BN differ sharply from those of graphite. Thus, BN has a low
electric conductivity, is colorless. and its diamagnetic susceptibility is small
while that of graphite is abnormally large. These properties are in agreement with a single bond model”.
Cubic BN has a diamond-like structure, each B atom is bonded to four N
atoms at a bond distance of 1.56 A. The enthalpy of the transformation
from hexagonal to cubic BN is 0.84 kcal mol-’ [131], and the energy of
atomization is AEZ (BN, cub) = AEf,(BN, hex) - 0.8 kcal mol-’ =
306.5 kcal mo1-l. By assuming that the energy of atomization is equal to
the energy required to break three single covalent bonds and one dative
bond, AE:, = 3 E(B-N) + E(N + B). and that E ( N B) = AH:,.
(H,N
BH,) = 31.1 kcal mol-I, we obtain E(B-N) = 92 kcdl mol-’.
-
-
1001
may be described as having equal covalent and dative bonding contributions:
normal : dative
l:o
R
1002
Me
R=Me
3:1
Scheme 5. Canonical structures of [AI(NMe,),],
The results of molecular orbital bond index calculations
on [AI(NH,),], and similar compounds have just been publ i ~ h e d . [ ' ~ ~ ~authors
T h e note that the bond index of the AI-N
terminal bond ist nearly twice the bond index of the bridging
AI-N bond and conclude: "This indicates a greater strength
for the terminal bonds [. . .I than previously estimated, the
difference arising because the capacity of terminal ligands
with lone pair electrons to indulge in dative %-bonding has
been underestimated in earlier work". As will be clear from
the preceding discussion, we prefer to regard the terminal
bonds as single covalent bonds, and the bridging bonds as
less than that, viz. as half covalent and half dative. Using the
terminal AI-N distance as a reference value for a covalent
AI-N bond and the mean of the N +A1 bond distances in
Me,NAlH, and Me,NAIMe, (2.08 A) as a reference value
for a dative bond distance, we calculate a mean value of
1.95 A, in reasonable agreement with the observed value of
I .97 A. A description of each of the four bonds as half covalent and half dative is consistent with the expected fragmentation behavior: two AI-N bonds are expected to break heterolytically to yield the monomeric amides, the next two to
break homolytically to yield (Me,N),AI' and "Me,.
In the following we shall briefly discuss the bond distances
in some polynuclear AIN compounds (see Fig. 7). In order to
minimize inductive effects we confine ourselves to species
with hydrogen and alkyl substituents.
The AI-N
bonds
in
(H,AINMe,),[1321 and
(Me,A1NMe,),['341 may also be described as half covalent
and half dative. The calculated mean (1.95 A) is in good
agreement with the observed values.
The AI-N bond distance in the cage compounds (HAINiPr)," 351 and (HAINnPr),['361 may be described in terms
of 2: 1 covalent and dative bonding contributions. From the
reference values we calculate a bond distance of 1.90 A,
again in good agreement with the observed values.
Finally, one may describe the bonds in crystalline
AIN[1371as 3:l resonances between covalent and dative
bonds. The corresponding calculated bond distance is
1.88 A. Again, a linear, two-parameter model has been able
to reproduce the AI-N bond distances in Figure 7 with a
deviation of 0.02 A or less.
This mode1 is not equally successful in all cases. Thus, the
0.04 A
Al-N bridge bonds in (CI,AINMe,), are 1.91
shorter than calculated from the reference bond distances.
The deviation may be interpreted as an inductive effect of the
electron withdrawing CI atoms. But adoption of the AI-N
and N +Al distances in 8 as reference values does not lead
to improvement, the calculated mean is 1.87 A, 0.04 A less
than the observed value.
=
2:1
R
= I-C,H,
R = n-C,H,,
kAI-N
= 1.91 A
1:l
R=Me
R=Me
0: 1
R=Me
R=Me
Fig. 7. The structure of some polynuclear AIN compounds and the ratio of
normal to dative contributions to the AI-N bonds.
But even in those cases where the model fails in a quantitative way, it remains true that bond distances generally
increase with increasing dative bonding contributions. The
molecular structures of 10-14 have all been determined by
(Me,AI), (p-PMe,),
(Me,AI),(p-SMe),
and (Me,AI),(p-CI),
(Me,AI), (p-OMe), 11 11401
12"411 (Me2Al)4(p-F)4 13"421
141i431
gas electron diffraction. In all these compounds we describe
the endocyclic bonds as half dative and half covalent. The
AI-P bond in 10 (2.43 A) is 0.10 8, shorter than in adducts
between AIMe, and phosphanes,['*. 1441 and 0.06 A longer
than the bond distance in solid aluminum phosphide, which
we would regard as 314 covalent and 1J4 dative; we are not
aware of any molecule containing what might be described
as a single covalent AI-P bond, but expect it to be about
2.32
A.
The AIL0 bond distance in 11 (I .85 A) is close to the mean
of the terminal covalent AI-0 distance in tetrameric aluand the dative 0 -+A1
minum trisisopropoxide (1.72
distance in Me,O +AIMe, (2.01(2) A). The A1-S bond distance in 12 (2.37 A) is about 0.18 A shorter than the dative
S +A1 distance in Me,S +A1Me3,[841while the AI-F distance in 13 (1.81 A) is 0.18 A greater than the normal AI-F
distance in gaseous monomeric AIF, .11451
Finally, the AI-CI
Angen. Chem. Int. Ed. Engl. 28 (1989) 992-1007
bond distance in 14 (2.30 A) is 0.24 8,longer than the normal
bond distance in gaseous, monomeric AlCI, .[461
Before going on to other metals we pause to take note of
the dinuclear complex 15. containing a five-coordinate A1
atom,"461 and the cationic complex 1611471prepared with
[AICI,]@ as counterion.
covalent contribution to the axial Al(ltN(4) bond at the
expense of the equatorial AI(IFN(1) bond.
The mean AI-C distance in the cationic complex 16 is not
significantly different from the AI-C bond distances in free
AlMe, or Me,NAlMe,, but the dative N +A1 distances,
1.92(2) A, are 0.18 8, shorter than in the latter adduct. It
should be noted, however, that even this very short N +Al
bond distance is about 0.11 A greater than the reference
value for a covalent Al-N bond that we adopted above.
11. Three-coordinate Lithium Amides
15
The structure of the adduct of tetramethylethylenediamine and a monomeric Li-amide is shown in Figure 8.['481
A ryl,
The five-coordinate Al(1) atom in 15 is trigonal bipyramidal, the four Al(lFN bond distances range from 1.83 to
2.14 8,.The AI-C and Al(ltN(3) bonds (1.83 A) are clearly
covalent, the Al(ltN(2) bond (2.14 A) clearly dative. As
expected the covalently bonded C and N(3) occupy equatorial positions, the datively bonded N(2) an axial position.
Al( 1) is bonded to the four-coordinate atom Al(2) through
two amido nitrogen atoms, N(l) and N(4). We have seen that
when two four-coordinate Al atoms are connected through
two amido bridges as in (Me,AINMe,),, covalent and dative
bonding contributions are delocalized to yield four equal
Al-N bonds at about 1.95 A. In 15 the bond distances in the
four-membered AI,N, ring range from 1.91 to 2.06 8,.
There are probably several reasons for this variation.
Firstly, as we have seen, bond distances to the acceptor atom
are expected to increase with increasing coordination number, the mean of the bond distances from N((1) and N(4) to
the five-coordinate A1 is about 0.10 8, greater than the mean
of the bond distances to the four-coordinate Al. Secondly, in
compounds of the main group elements with trigonal bipyramidal coordination the bonds to axial ligands tend to be a
few hundreths of an 8, unit longer than those to equatorial
ligands. Finally, we interpret the small but significant difference between the bond distances to the four-coordinate
Al(2) (1.935(2) 8, and 1.911(3) A) as an indication that covalent and dative bonding contributions are not completely
delocalized over the AI,N, ring: the shortest bond, Al(2)N(4). and consequently also the AI(1)-N(l) bond appear to
be slightly more than half covalent, the longer bonds Al(2)N( 1) and Al( 1 )-N(4) appear to be slightly more than half
dative. We further suggest that the reason for the incomplete
delocalization is that it would require energy to increase the
AI-C 1.96(2) 8,
3 CAlC 152"
3 NAIN 115'
16
Angrir. Chem. Ini. Ed. Engl. 28 (1989) 992-1007
N
normal :dative
,H
l:o
1.90 8,
2.15 A
0: 1
v, p /Y2 0l 4 A
1:1
",
u
Et,O+ Li,/ ' .,LI +OEt,
N,
H Aryl
1:l
RN-
/I
Li YNR
J
"
O
9
/I
LI
2.16A
01
1'jSi-!TNR
RN-
Si,
Aryl = 2,4,6-tBu3C,H,; R = tBu
Fig. 8. Bond distances in some three-coordinate lithium amides and the ratio
of normal to dative contributions to the Li-N bonds.
Here again we are able to compare dative and normal bond
distances between the same atom pair in the same molecular
species: Li-N = 1.90 A and N + Li = 2.15 A.
In [Et,OLi(p-NHAryl),], (Fig. 8, center) the Li atom is
datively bonded to the 0 atom, while each of the four endocyclic bonds must be regarded as half dative and half norma1.['491The observed Li-N distance is in good agreement
with the mean of the pure normal and dative bond distances
in the adduct above shown at the top of Figure 8, 2.03 A.
Finally, we turn our attention to the cubic complex at the
bottom of Figure 8.[lSo1The Si atoms are seen to be fourcoordinate, and all Si-N distances are of the magnitude
expected for covalent bonds. Since the Si-N bonds in the
basal plane and the terminal N-C bonds are covalent, it
follows that the perpendicular N-Li bonds are dative as
indicated. If N is assumed to be capable of forming three
covalent and one dative bond, and Li to be capable of forming one normal and two (or more) dative bonds, it follows
that since the Li-N bonds in the upper face are all similar,
each must be half normal and half dative. The observed
mean value, 2.09 A, is in moderate agreement with the estimated value of 2.03 A.
1003
12. Two Associated Si Amides
Me,NSiH,CI is monomeric in the gas phase, but dimeric
in the s ~ l i d . ~ The
’ ~ ’ monomeric
~
units in the dimer 17 are
held together by two dative N + Si bonds. The coordination
around the two Si atoms is trigonal bipyramidal with the
datively bonded N atom and the CI atom in apical positions.
The Si-N and N + Si bond distances are comparable to
those in the mononuclear adducts 4 and 5 discussed in Section 7.
covalent and dative bonding contributions; the mean of the
Zn-0 and 0 + Zn bond distances in 19 is 1.99 A, in good
agreement with the observed value.
19, Aryl = 2,4,6-tBu3C,H,
Tetrameric MeZnOMe has the cubane structure 20.”’
The external Zn-C bond distance corresponds to a covalent
17
We have seen that when two NR, groups bridge two fourcoordinate Al or two three-coordinate Li atoms, covalent
and dative bonding is delocalized to give four equal bond
distances. We suggest that this does not take place in the
present compound because it requires energy to increase the
covalent bonding contribution to the axial Si-N bond at the
expense of the equatorial bond.
Replacement of the C1 atom by a hydrogen atom leads to
a radical change of the solid state structure: Me,NSiH, is
pentameric in the solid phase as indicated in 18.[1521
bond. If it is assumed that Zn is capable of forming one more
covalent bond and two (or more) dative bonds, it follows
that each Zn-0 bond in the cube is 1/3 covalent and 2/3
dative. The observed bond distance in 20 is indeed longer
than the 1 : 1 bond in solid ZnO and shorter than the pure
dative bond in 19, but it is significantly longer than the
appropriate mean of the bond distances in 19 (2.02 A), perhaps as an effect of the presence of an electron-releasing
methyl group.
14. “Tetrahedral Covalent Radii”
<N-Si-N,
H,C” I
I
CH, H
i ”CH,
CH,
18
The covalently bonded N atom has moved into the apical
position formerly occupied by the CI atom: formation of
dimers is prevented by the N-Si-N angle of 180”, the degree
of association is now determined by the valence angle at N.
The Si-N bond distances in the ten-membered ring of 18
are indistinguishable. Each of them may be described in
terms of equal covalent and dative bonding contributions,
though the observed bond distance is somewhat larger than
the average of the localized Si-N and N + Si distances in 17
(1.93 A). We suggest that delocalization is stabilizing in the
case of 18 because both N atoms occupy apical positions.
13. Some Zinc-Oxygen Compounds
The Zn atom in the complex 19 formed from Zn(OAryl),
and two moles of tetrahydrofuran is covalently bonded to
two 0 atoms and datively bonded to two others.[1531The
difference between the two bond distances is 0.19 A.
Solid ZnO has a wurzite structure, each Zn atom is surrounded by four 0 atoms at an average distance of
1.98
The bonding may be described in terms of equal
3 004
Working with the evidence available at the time when The
Nature of the Chemical Bond“] was written, Pattling clearly
found no reason to distinguish between covalent and dative
bonds. He described bonding in solid ZnO, AlN or cubic BN
as “covalent bonds with some ionic character”. N and 0
atoms were assigned tetrahedral covalent radii equal to the
normal covalent radii determined from the observed bond
distances in amines, ethers and related compounds. The “tetrahedral covalent radii” of Zn and the other Group 12 metals were then adjusted to reproduce the observed interatomic distances in oxides and other binary compounds with
Group 16 elements. Similarly, the radii of Al and the other
Group 13 elements were adjusted to reproduce the observed
interatomic distance in nitrides and other binary compounds
with group 15 elements.
The observed interatomic distances that we have adopted
as reference values for pure, single covalent bonds to fourcoordinate metal atoms are consistently shorter than the
estimates obtained by adding the “tetrahedral covalent radii” for the two atoms: for A1-N 1.81 A versus 1.91 A, for
AIL0 1.71 A versus 1.92 A, and for Zn-0 1.89 A versus
1.97 A. For B-N we have adopted 1.47 A as a reference
value for a single covalent bond to trigonal B; the sum of the
“tetrahedral covalent radii“ is 1.58 A.
Angew. Chem. Int. Ed. Engl. 28 (1989) 992-1007
The reason for the discrepancy is clear. In our terminology
the bonds in solid ZnO, AIN or cubic BN are not pure
covalent bonds, but bonds of intermediate type with significant dative bonding contributions. The "tetrahedral covalent radii" are therefore incapable of reproducing either of
the bond distances in species like 16, which contain what
according to our terminology are pure covalent and dative
bonds.
We have quoted a few instances where this discrepancy has
led investigators to the conclusion that the covalent bonds
are shortened by additional .n-bonding; there are more instances in the literature.
If the distinction between covalent and dative bonds is
made, there arises the need for a new set of covalent radii
based on observed pure, single covalent bond distances. In
view of the great sensitivity of dative bond distances to inductive effects, it would hardly be profitable to attempt to
establish a set of dative bond radii.
15. Conclusion
In this review we have defined a chemical bond as normal
if minimum-energy rupture proceeds homolytically and as
dative if minimum-energy rupture proceeds heterolytically.
The term normal covers covalent and ionic bonds as well
as bonds of intermediate polarity. We have compared the
physicochemical properties of chemical bonds between the
same atom pair DM where D is a donor element in groups
15, 16 or 17 and M is a main group metal: bonds to transition metals have not been considered. We find that:
1. The strength of a dative, or donor-acceptor, bond is very
sensitive to inductive effects, particularly at the acceptor
atom. Thus, while SnMe, does not appear to interact with
electron donors such as amines or pyridine, the N + Sn bond
rupture enthalpy of pyr + SnMe,C1 is 11 kcal mol-' and of
Me,N + Sn@Me, 46 kcal mol-'. The variation in bond
strength is clearly reflected in bond distances. Thus, while the
N --+ Si bond distances in neutral (five-coordinate) Si" adducts range from 2.70 to 1.95 8, with increasing electronegativity of the substituents, the N + Si bond distance in
pyr + Si@Me,is 1.86 A.
2. The available evidence for metals in groups 12,13 and 14
suggests that the strength of a dative bond rarely exceeds half
the strength of an isoelectronic normal bond or a normal
bond between the same atom pair.
Again the difference is reflected in bond distances : covalent Si-N distances to four- or five-coordinate Si" range
from 1.70 to 1.80 A, dative N + Si distances range from 1.95
to 2.70 A. Normal and dative AIN bond distances in the
same molecular unit 8 are 1.77 and 1.96 A, respectively. Normal and dative LiN bond distances in the adduct at the top
of Figure 8 are 1.90 and 2.15 A, respectively, and the ZnO
bond distances in 19 are 1.89 and 2.08 A.
3. The coordination geometry of the acceptor atom may be
predicted from the VSEPR model with the additional postulate that the spatial requirements of a (partially) "accepted"
bond pair are less than the spacial requirements of a normal
bond pair (NBP > ABP). This means that in five-coordinate
Angew. Chem. I n f . Ed. Engl. 28 (1989) 992-1007
trigonal pyramidal adducts the electron donors will occupy
the axial positions.
The VSEPR model leads to the prediction that bonds to
the acceptor atom will be elongated and the angles between
them reduced on coordination of (another) electron donor.
Both effects will be reinforced by steric repulsion between
donor molecule and acceptor atom substituents.
4. In some molecular species bond distances are found to be
equal, though no reasonable Lewis structure in which all are
covalent or all dative can be found. Bonding in such species
is conveniently described in terms of canonical forms containing localized covalent and dative bonds. Delocalization of
bonding (resonance) yields bonds that are intermediate between pure covalent and dative bonds. The covalent or dative contributions to a particular bond in the real molecule
may be determined by counting the number of canonical
forms in which it is covalent or dative. The lengths of such
bonds are generally found to be intermediate between the
lengths of a pure covalent and a pure dative bond and to
increase with increasing dative bonding contributions. In
favorable cases bond distances may be predicted to within
two or three hundreths of an 8, unit by calculating the appropriate mean of reference values for pure normal and pure
dative bond distances.
This approach is particularly useful for describing bonding in symmetric M,(p'-NR,), M,(p'-OR) or M,(p2-X)
fragments; each M-N/O/X bond is half normal and half
dative. Similarly, each M-N bond in a symmetric M,(p-NR)
fragment is 2/3 normal and 1/3 dative, and each M-O/X
bond in symmetric M,(p3-OR) or M,(p3-X) fragments are
1/3 normal and 2/3 dative.
We are grateful to the Norwegian Research Council for Science and the Humpities (NAVF) for financial support. We
thank. Anne-Lise Agren for typing - and retyping - the manuscript, and Snefrid Gundersen for preparing the figures.
Finally I am grateful to Professors M . E Lappert, P . Paetzold
and P . von R. Schleyer and two anonymous referees for constructive criticism and helpful comments.
Received: January 26, 1989;
revised: February 20, 1989 [A 729 IE]
German version: Angew. Chem. 101 (1989) 1017.
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