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Journal ArticleDOI

Lithium bonding interaction in H 2 CY⋯LiF (Y=O,S)complexes: A theoretical probe

15 Sep 1997-Journal of Chemical Physics (American Institute of Physics)-Vol. 107, Iss: 11, pp 4329-4336
TL;DR: In this paper, the effect of correlation on complex binding energies has been studied via single point MP2 (full) calculations done on 6-31++G(d,p) geometry.
Abstract: Ab initio calculations at 6-31++G(d,p) level have been done on H2CY⋯LiF (Y=O,S) complexes choosing ten possible orientations in each complex. The effect of correlation on complex binding energies has been studied via single point MP2 (full) calculations done on 6-31++G(d,p) geometry. Binding energies have been corrected for basis set superposition error. Frequency calculations confirm that H2CO⋯LiF and H2CS⋯LiF complexes have three and two stable forms, respectively. The most stable form in each complex has been found to have a strong lithium bonding interaction and a secondary hydrogen bonding interaction. NBO analysis has revealed that in this form oxygen donates nσ lone pair while sulfur donates its nπ lone pair. In yet another stable form of these complexes, mixed donation of π and nσ electrons have been observed.

Summary (1 min read)

1. Introduction

  • During the last 900 000 years, the Earth’s climate has been characterized by a strong cyclicity of 100 000 years leading to the alternation of short interglacial periods, which developed within a more or less severe glacial climate.
  • These icebergs were carried by the surficial oceanic currents, drifted and melted into the North Atlantic Ocean, north of 40◦N [3–5].
  • At this point, neither mechanism is widely accepted.
  • During this period, the ice sheets were at a size intermediate between full glacial and full interglacial conditions, and the insolation forcing did not experience large variations.

2. Chronostratigraphical constraints

  • Identifying the same event in several cores raised from a large oceanic basin requires a reliable chronostratigraphy.
  • For H4, the error bars associated with the 14C dates are up to several hundred or thousands of years and calendar ages were computed using the polynomial equation of Bard [1].
  • To take into account the trend of global ice volume changes recorded during these events, the δ18O amplitude has been corrected by −0.08 for H1 (estimated following [2]) and +0.06 for H4 (estimated following [39]).

3. Results: iceberg origin and transport

  • The amplitude of δ18O variations associated with H4 shows a pronounced latitudinal structure in the meltwater extension.
  • The δ18O amplitude is high in both the Norwegian Sea and the North Atlantic Ocean north of 40◦N, but decreased from the west to the east.
  • These data suggest a major iceberg discharge originating from both North American and North European ice sheets.
  • Since planktonic Fig. 3. Changements des conditions hydrologiques de surface associés à l’événement de Heinrich 1, reconstruits à partir de l’amplitude isotopique (δ18O avant H1 – δ18O pendant H1).
  • The extension of the meltwater belt is not as latitudinal as in H4, and values <1 are present in the eastern Atlantic basin.

4. Discussion

  • The impact of both Heinrich events on the surface hydrology of the North Atlantic Ocean has been very different.
  • Several hypotheses can be proposed: cooling intensity: if the cooling during H4 was higher than during H1, the amplitude of the isotopic anomaly would be smaller, even if the magnitude of iceberg discharges was roughly the same.
  • Some similar modelling experiments should be performed for H1 to better constrain the characteristics of this event.
  • The authors used also SST reconstructions whenever possible.
  • The meltwater area was not fully developed from west to east in the 40–50◦N latitudinal band and the Gulf Stream and North Atlantic Drift were not stopped.

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University of South Carolina University of South Carolina
Scholar Commons Scholar Commons
Faculty Publications Chemical Engineering, Department of
1997
Lithium bonding interaction in Lithium bonding interaction in
HH
22
CYLiF CY LiF
(Y(Y
=O,S=O,S
))
complexes: A complexes: A
theoretical probe theoretical probe
Salai Cheettu Ammal
University of South Carolina - Columbia
, ammal@cec.sc.edu
P. Venuvanalingam
S. Pal
Follow this and additional works at: https://scholarcommons.sc.edu/eche_facpub
Part of the Biological and Chemical Physics Commons
Publication Info Publication Info
Published in
The Journal of Chemical Physics
, Volume 107, Issue 11, 1997, pages 4329-4336.
Copyright 1997 American Institute of Physics. This article may be downloaded for personal use only. Any
other use requires prior permission of the author and the American Institute of Physics.
The following article appeared in
Ammal, S. S. C., Venuvanalingam, P., & Pal, S. (1997). Lithium bonding interaction in H2CYLiF (Y=O,S)
complexes: A theoretical probe.
The Journal of Chemical Physics
, 107(11), 4329-4336. http://dx.doi.org/
10.1063/1.474773
and may be found at
http://scitation.aip.org/content/aip/journal/jcp/107/11/10.1063/1.474773
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Lithium bonding interaction in H 2 CY⋯LiF ( Y = O,S ) complexes: A theoretical probe
S. Salai Cheettu Ammal, P. Venuvanalingam, and Sourav Pal
Citation: The Journal of Chemical Physics 107, 4329 (1997); doi: 10.1063/1.474773
View online: http://dx.doi.org/10.1063/1.474773
View Table of Contents: http://scitation.aip.org/content/aip/journal/jcp/107/11?ver=pdfcov
Published by the AIP Publishing
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Lithium bonding interaction in H
2
CY•••LiF (Y5 O,S) complexes:
A theoretical probe
S. Salai Cheettu Ammal and P. Venuvanalingam
Department of Chemistry, Bharathidasan University, Tiruchirappalli620 024, India
Sourav Pal
Physical Chemistry Division, National Chemical Laboratory, Pune411 008, India
~Received 6 February 1997; accepted 11 June 1997!
Ab initio calculations at 6-3111G~d,p! level have been done on H
2
CY•••LiF ~Y5O,S! complexes
choosing ten possible orientations in each complex. The effect of correlation on complex binding
energies has been studied via single point MP2 ~full! calculations done on 6-3111G(d,p)
geometry. Binding energies have been corrected for basis set superposition error. Frequency
calculations confirm that H
2
CO•••LiF and H
2
CS•••LiF complexes have three and two stable forms,
respectively. The most stable form in each complex has been found to have a strong lithium bonding
interaction and a secondary hydrogen bonding interaction. NBO analysis has revealed that in this
form oxygen donates n
s
lone pair while sulfur donates its n
p
lone pair. In yet another stable form
of these complexes, mixed donation of
p
and n
s
electrons have been observed. © 1997 American
Institute of Physics. @S0021-9606~97!01035-0#
INTRODUCTION
The weak interaction between closed-shell molecules
plays a vital role in an enormous variety of chemical, physi-
cal, and biological phenomena.
1–4
The basic understanding
of such weak but central interactions is necessary to enable
design and manipulation of molecular systems that depend
on noncovalent binding. There are many experimental tech-
niques available to study the intermolecular complexes in
condensed phases, which can provide information about the
structure and energetics of the complexes.
5
But the informa-
tion obtained is, of course, affected by cooperative phenom-
ena operating in condensed phases. The detailed knowledge
about the specific bimolecular interaction can be obtained
only if the study is carried out in gas phase or in inert sol-
vent. Experimental methods may be powerful for gas phase
study of the complexes, but none of them provides quantita-
tively accurate results on structural information as well as
interaction energies. This makes the quantum chemical
methods an important tool to study intermolecular
interactions.
6
Intermolecular interactions can be categorized mainly
into three classes depending on their nature and strength.
They are H-bonding, chargetransfer, and van der Waals
interactions. Among them, hydrogen bonded systems have
been well studied because of their significant role in chemi-
cal and biological interactions.
7
Correspondingly, ab initio
calculations on them are becoming increasingly common in
the literature.
8
Analogous to hydrogen, lithium can also par-
ticipate in a three center interaction known as lithium bond-
ing. While hydrogen bonding has been so widely and thor-
oughly investigated, reports on lithium bonding are very
rare.
9–12
As the experimental studies on the lithium bond is
very meagre, most of our current knowledge relating to
lithium bonding has been derived from theoretical work.
13,14
Lithium bonding was speculated as early as in 1959 by
Shigorin,
9
and later by West and Glaze
10
and Brown and
co-workers.
11
While Ault and Pimental
12
were the first to
provide experimental proof for the existence of a lithium
bond in H
3
N•••LiX ~X5Cl,Br! complexes, Kollman and
co-workers
13
were the first to theoretically investigate the
properties of lithium bonds using ab initio self-consistent-
field ~SCF! calculations. Subsequently there were few theo-
retical reports published on lithium bonded dimers.
14
Though
lithium bonding closely resembles hydrogen bonding in sev-
eral ways, there are many differences between these two
types of bonding.
15
Notable among the differences is that
lithium bonding is stronger than hydrogen bonding. Another
important difference between hydrogen and lithium bonds is
that the electrostatic interaction dominates over charge
transfer interaction in the lithium bonding, whereas charge
transfer interaction also contributes significantly in the case
of hydrogen bonding.
16
It is this feature that makes the com-
parative study of hydrogen and lithium bonding interactions
interesting. The present study focuses on intermolecular in-
teractions of lithium fluoride, a powerful lithium donor, with
the prototype bases formaldehyde and thioformaldehyde,
through theoretical calculations.
Formaldehyde, which is known as the most common
base to form a hydrogen bond with a proton donor in many
organic and biological systems, has different donor sites. It
can act as a lone pair (n
o
) donor and a
p
~.C5O! donor.
Further, the hydrogens of formaldehyde can form weak hy-
drogen bonds through CH protons with proton acceptors
such as halogen—especially fluorine—oxygen, sulfur, etc.
Hydrogen bonded complexes of formaldehyde have been the
subject of several theoretical
17
and experimental
18
investiga-
tions. Lithium halides and lithium hydrides can also form
complexes with formaldehyde in a similar way and yet stud-
ies of lithium bonded complexes, even with molecules hav-
ing first row atoms, are rare. Lithium fluoride is chosen be-
cause the complex can simultaneously have both lithium and
hydrogen bonding interactions and hence can have remark-
4329J. Chem. Phys. 107 (11), 15 September 1997 0021-9606/97/107(11)/4329/8/$10.00 © 1997 American Institute of Physics
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able stability. Further, LiF can exist almost as Li
1
F
2
and
there is a possibility that the ionic species Li
1
and F
2
can
interact at different sites of formaldehyde and form ionic
molecular complexes. Singh and co-workers
19
have studied
the complex of NH
3
with LiCN, and proved that in the
ground state, minimum energy geometry of the complex Li
1
binds with nitrogen and CN
2
binds with the hydrogen of
NH
3
. Besides formaldehyde, its thioanalogue thioformalde-
hyde is chosen to compare the directionality of lithium bonds
to sulfur and oxygen. It is known
20
that sulfur forms slightly
weaker hydrogen bonds with more ‘‘perpendicular’’ angles
than oxygen, and this is believed to result from different
hybridizations of valence orbitals in oxygen and sulfur. Re-
cent calculations
21
show that the reason for the different geo-
metrical preferences is not as simple as has been explained
above. The above report further reveals that the charge
charge interaction dominates in the O•••H bond that prefers
the linear orientation, while the chargequadrupole interac-
tion plays a significant role in the S•••H interaction that pre-
fers a perpendicular orientation. Therefore the hydrogen
bond formed between sulfur and hydrogen is fundamentally
different from the O•••H bond. It is therefore interesting to
look at the geometry and the factors responsible for the geo-
metric preferences of the lithium bonded complexes with
donors formaldehyde and thioformaldehyde. The results of
the above complexes will be compared with those of
H
2
CY•••HF complexes available in the literature
21,22
with a
view to observing the effect of the Li bond in complexation.
COMPUTATIONAL DETAILS
Different possible geometries considered for the com-
plexes H
2
CY•••LiF where Y5O,S are shown in Fig. 1. Fol-
lowing are the details of fixation of various geometric param-
eters in optimizing the selected structures of the above
complexes. Structure I has been optimized with C
2
v
symme-
try constraint such that the LiF molecule is placed on the C
2
axis. C
2
v
symmetry has been relaxed to C
s
symmetry in
structure II by allowing all the bond lengths and bond angles
to vary and fixing the torsion angles 5-1-2-3 and 6-5-1-2 at
0°. Structure III is planar and therefore has C
s
symmetry. In
this structure LiF is positioned on H
2
CY such that both
Y•••Li and F•••H bonds are present. In structure IV, LiF has
been constrained to move within the bisecting vertical plane
of H
2
CY with Li anchored on the Y atom. In such an orien-
tation the torsion angle of H–C–Y–Li is always 90°. Struc-
tures V and VI have been optimized with C
s
and C
2
v
sym-
metry, respectively, and in that V and VI have linear and
bifurcated hydrogen bonds as shown in Fig. 1. In structure
VII, electron donation from .C5O
p
-bond is considered.
The LiF molecule is placed vertically at the midpoint of the
p
-bond with the Li atom pointing downward towards the
p
-bond and the structure has been fully optimized. LiF can
also exist in ionic form as Li
1
F
2
, and in that the ionic spe-
cies Li
1
and F
2
can bind with H
2
CY at different sites. Vari-
ous possibilities of this kind have been tried out with struc-
tures VIII, IX, and X presented in Fig. 1. In all these
structures Li
1
binds with oxygen/sulfur atom and F
2
forms
bifurcated ~VIII!, linear ~cis!~IX!, and linear ~trans!~X!hy-
drogen bonds with hydrogen atoms of H
2
CY. C
s
symmetry
has been maintained in all the three cases. The optimized
structures of II and III have then been used to scan through
the potential energy surface. Single point energies have been
calculated with PES scan for change in the torsion angle
3-2-1-5 ~structures II and III! from to 90° and in steps of
10°.
The geometrical parameters of the complexes have been
optimized using the supermolecule approach at SCF level.
Pople’s split-valence double-zeta 6-31 G basis set augmented
by one set of d-polarization functions on heavy atoms
and p-polarization function on hydrogens and also sp-
~heavy atoms! and s- ~hydrogens! diffuse functions
@
6-3111G(d,p)]
23
is used in the calculations. Addition-
ally, single point energy calculations on the SCF optimized
geometries have been carried out at the MP2 ~full! level in
order to include electron correlation correction to the inter-
action energy. The frequency calculations have been carried
out to confirm the nature of the stationary points obtained.
The interaction energies have been corrected for basis set
superposition error ~BSSE! using the BoysBernardi
24
coun-
terpoise scheme and applying a modification
25
that takes into
FIG. 1. Different possible geometries of H
2
CY•••LiF ~Y5O,S! complexes.
4330
Ammal, Venuvanalingam, and Pal: H
2
CY•••LiF (Y5O,S) complexes
J. Chem. Phys., Vol. 107, No. 11, 15 September 1997
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consideration the geometrical relaxation of the monomers
upon complexation. Natural bond orbital ~NBO! analysis
26
on the stable forms III and IV of the complexes have been
carried out to examine the nature of the interaction at the
orbital level. All the calculations have been carried out using
the
HONDO7 program
27
implemented on a Digital Dec Alpha
system.
RESULTS AND DISCUSSION
Lithium fluoride is both a lithium donor and a proton/
lithium acceptor. Similar to LiCN, it can exist as an ion pair
and therefore form ion pair molecular complexes and ionic
molecular complexes. Formaldehyde and thioformaldehyde
are proton/lithium acceptors and can donate their
s
or
p
type
lone pair or
p
bond pair to form ~Lewis acid base! com-
plexes. Further, the above bases have two protons and
through them they can form either linear or bifurcated
H-bonds. This is more so in the complex stage where
formaldehyde/thiofomaldehyde is sufficiently polarized to
donate its protons. All these possibilities are considered here
and are shown in Fig. 1. There are 20 total species—10 for
formaldehyde and another 10 for thioformaldehyde—the ge-
ometries of which have been optimized. The binding ener-
gies have been calculated at HartreeFock level with basis
functions of double-zeta quality augmented by diffuse and
polarization functions on both heavy and hydrogen atoms. In
H
2
CO•••LiF complex optimizations with starting geometries,
VIII, IX, and X converged on structure III, and this shows
that ionic complexes in these cases are absent. And the com-
plex with initial geometry VII converged on IV. Thus struc-
tures IVI alone represent stationary points on the PES, and
in them I, V, and VI have turned out to be first order saddles.
Therefore only structures II, III, and IV represent equilibrium
geometries. Among them, structure III appears to be the most
stable. In the H
2
CS•••LiF complex again, optimization with
starting structure VIII and IX converged on III and that, with
structure X, unlike the formaldehyde case, converged to a
geometry with a markedly low value of the F
2
•••H distance
~0.93 Å!, indicating a proton transfer. This shows that there
are no ionic complexes present and, unlike the H
2
CO•••LiF
complex, optimizations with initial geometries II and VII
converged on III and IV, respectively. Therefore there are
only five stationary points corresponding to structures I, III,
IV, V, and VI. In them structures I and V have turned out to
be first order saddles and VI a second order saddle. Fre-
quency computations show that structures III and IV are
stable equilibrium structures and between them III appears to
be more stable. Beside these, PES scans have been made
searching for possible minimum with LiF interacting with
these bases from the molecular plane by changing H–C–
YLi torsion from to 90° in steps of 10°, and no stable
geometry other than IV has been found in the process. Inter-
action energies, BSSE, etc., have therefore been calculated
for the first six structures alone.
Interaction energies, number of imaginary frequencies
obtained for each structure, BSSE, and counterpoise cor-
rected interaction energies are presented in Table I. Effect of
correlation on complex binding energy has been studied only
for the stable forms II, III, and IV in the case of H
2
CO•••LiF
complex and III and IV of H
2
CS•••LiF complex using MP2
~full! calculation on 6-3111G(d,p) optimized geometry,
and the results are presented in Table II. Selected geometric
parameters of the monomers and complexes ~III and IV! are
presented in Table III. NBO analysis and orbital energy cor-
relation have been done only for III and IV. Quantum of
charge transfer q
CT
, occupancies of the orbitals participating
in the donoracceptor interaction, and the second order en-
ergy lowering due to the interaction of the donor and accep-
tor orbitals DE
(2)
obtained from NBO analysis, are collected
in Table IV. Table V presents the orbital energies of mono-
mers and complexes and, for obvious reasons, the energies of
only a few top lying acceptor and donor orbitals have been
presented.
TABLE I. Interaction energies DE
b
, BSSE, counterpoise corrected interac-
tion energies DE
b
cp
~kcal/mol! and number of irnaginary frequencies (n
i
) for
the complexes of LiF with H
2
CO and H
2
CS calculated at HF/6-31
11G(d,p) level.
Structure
H
2
CO•••LiF H
2
CS•••LiF
DE
b
BSSE
DE
b
cp
n
i
DE
b
BSSE
DE
b
cp
n
i
I 18.75 1.88 16.87 1 4.46 1.16 3.30 1
II 18.88 1.71 17.16 0 ••• ••• ••• •••
III 20.87 1.00 19.87 0 16.14 0.53 15.61 0
IV 17.17 0.78 16.39 0 9.40 0.81 8.59 0
V 5.59 0.30 5.29 1 6.06 0.28 5.78 1
VI 5.87 0.33 5.54 1 5.07 0.30 4.77 2
H
2
CO•••HF
H
2
CS•••HF
4-31G
a
10.4 ••• ••• ••• 6.9 ••• ••• •••
a
Reference 22.
TABLE II. MP2/6-3111G~d,p! interaction energies DE
b
, BSSE, counterpoise corrected interaction energies
DE
b
cp
for the complexes of LiF with H
2
CO and H
2
CS ~kcal/mol!.
Structure
H
2
CO•••LiF H
2
CS•••LiF
DE
b
BSSE
DE
b
cp
DE
b
BSSE
DE
b
cp
II 20.63 5.62 15.01 ••• ••• •••
III 21.57 3.98 17.60 17.87 4.22 13.65
IV 17.44 3.61 13.83 9.85 4.45 5.40
H
2
CO•••HF
H
2
CS•••HF
MP2/6-31111G(d,p)
a
7.72 1.12 6.60 6.33 1.74 4.59
a
Reference 21.
4331
Ammal, Venuvanalingam, and Pal: H
2
CY•••LiF (Y5O,S) complexes
J. Chem. Phys., Vol. 107, No. 11, 15 September 1997
This article is copyrighted as indicated in the article. Reuse of AIP content is subject to the terms at: http://scitation.aip.org/termsconditions. Downloaded to IP:
129.252.69.176 On: Tue, 12 May 2015 20:09:24

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References
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TL;DR: In this paper, a direct difference method for the computation of molecular interactions has been based on a bivariational transcorrelated treatment, together with special methods for the balancing of other errors.
Abstract: A new direct difference method for the computation of molecular interactions has been based on a bivariational transcorrelated treatment, together with special methods for the balancing of other errors. It appears that these new features can give a strong reduction in the error of the interaction energy, and they seem to be particularly suitable for computations in the important region near the minimum energy. It has been generally accepted that this problem is dominated by unresolved difficulties and the relation of the new methods to these apparent difficulties is analysed here.

19,483 citations

Journal ArticleDOI
TL;DR: In this article, the 631G* and 6 31G* basis sets were extended through the second-row of the periodic table and the Hartree-Fock wave functions were used to obtain the equilibrium geometries for one-heavy-atom hydrides.
Abstract: The 6‐31G* and 6‐31G** basis sets previously introduced for first‐row atoms have been extended through the second‐row of the periodic table. Equilibrium geometries for one‐heavy‐atom hydrides calculated for the two‐basis sets and using Hartree–Fock wave functions are in good agreement both with each other and with the experimental data. HF/6‐31G* structures, obtained for two‐heavy‐atom hydrides and for a variety of hypervalent second‐row molecules, are also in excellent accord with experimental equilibrium geometries. No large deviations between calculated and experimental single bond lengths have been noted, in contrast to previous work on analogous first‐row compounds, where limiting Hartree–Fock distances were in error by up to a tenth of an angstrom. Equilibrium geometries calculated at the HF/6‐31G level are consistently in better agreement with the experimental data than are those previously obtained using the simple split‐valance 3‐21G basis set for both normal‐ and hypervalent compounds. Normal‐mode vibrational frequencies derived from 6‐31G* level calculations are consistently larger than the corresponding experimental values, typically by 10%–15%; they are of much more uniform quality than those obtained from the 3‐21G basis set. Hydrogenation energies calculated for normal‐ and hypervalent compounds are in moderate accord with experimental data, although in some instances large errors appear. Calculated energies relating to the stabilities of single and multiple bonds are in much better accord with the experimental energy differences.

6,870 citations

Journal ArticleDOI
TL;DR: In this paper, a method for extracting a unique set of atomic hybrids and bond orbitals for a given molecule, thereby constructing its Lewis structure in an a priori manner, is described.
Abstract: From the information contained in the (exact or approximate) first-order density matrix, we describe a method for extracting a unique set of atomic hybrids and bond orbitals for a given molecule, thereby constructing its “Lewis structure” in an a priori manner. These natural hybrids are optimal in a certain sense, are efficiently computed, and seem to agree well with chemical intuition (as summarized, for example, in Bent’s Rule) and with hybrids obtained by other procedures. Using simple INDO-SCF-MO wave functions, we give applications of the natural hybrid orbital analysis to molecules exhibiting a variety of bonding features, including lone pairs, multiple bonds, strained rings, and “bent bonds”, multiple resonance structures, hydrogen bonds, and three-center bonds. Three examples are described in greater detail: (i) “orbital following” during ammonia umbrella inversion, (ii) the dimerization of water molecules, and (iii) the hydrogen-bridged bonds of diborane.

4,338 citations

Book
01 Jan 1960

3,368 citations

Journal ArticleDOI
TL;DR: Developments in molecular and supramolecular design and engineering open perspectives towards the realization of molecular photonic, electronic, and ionic devices that would perform highly selective recognition, reaction, and transfer operations for signal and information processing at the molecular level.
Abstract: Supramolecular chemistry is the chemistry of the intermolecular bond, covering the structures and functions of the entities formed by association of two or more chemical species. Molecular recognition in the supermolecules formed by receptor-substrate binding rests on the principles of molecular complementarity, as found in spherical and tetrahedral recognition, linear recognition by coreceptors, metalloreceptors, amphiphilic receptors, and anion coordination. Supramolecular catalysis by receptors bearing reactive groups effects bond cleavage reactions as well as synthetic bond formation via cocatalysis. Lipophilic receptor molecules act as selective carriers for various substrates and make it possible to set up coupled transport processes linked to electron and proton gradients or to light. Whereas endoreceptors bind substrates in molecular cavities by convergent interactions, exoreceptors rely on interactions between the surfaces of the receptor and the substrate; thus new types of receptors, such as the metallonucleates, may be designed. In combination with polymolecular assemblies, receptors, carriers, and catalysts may lead to molecular and supramolecular devices, defined as structurally organized and functionally integrated chemical systems built on supramolecular architectures. Their recognition, transfer, and transformation features are analyzed specifically from the point of view of molecular devices that would operate via photons, electrons, or ions, thus defining fields of molecular photonics, electronics, and ionics. Introduction of photosensitive groups yields photoactive receptors for the design of light-conversion and charge-separation centers. Redox-active polyolefinic chains represent molecular wires for electron transfer through membranes. Tubular mesophases formed by stacking of suitable macrocyclic receptors may lead to ion channels. Molecular self-assembling occurs with acyclic ligands that form complexes of double-helical structure. Such developments in molecular and supramolecular design and engineering open perspectives towards the realization of molecular photonic, electronic, and ionic devices that would perform highly selective recognition, reaction, and transfer operations for signal and information processing at the molecular level.

3,124 citations

Frequently Asked Questions (16)
Q1. What contributions have the authors mentioned in the paper "Lithium bonding interaction in h2cy⋯lif (y=o,s) complexes: a theoretical probe" ?

This article may be downloaded for personal use only. Any other use requires prior permission of the author and the American Institute of Physics. The following article appeared in Ammal, S. S. C., Venuvanalingam, P., & Pal, S. ( 1997 ). 

which is known as the most common base to form a hydrogen bond with a proton donor in many organic and biological systems, has different donor sites. 

The net effect appears to stabilize the formaldehyde complex more than the thioformaldehyde complex, and the over-whelming change is brought out by a much stronger lithium bonding in the formaldehyde complex. 

LiF can exist almost as Li1F2 and there is a possibility that the ionic species Li1 and F2 can interact at different sites of formaldehyde and form ionic molecular complexes. 

The basic understanding of such weak but central interactions is necessary to enable design and manipulation of molecular systems that depend on noncovalent binding. 

C2v symmetry has been relaxed to Cs symmetry in structure II by allowing all the bond lengths and bond angles to vary and fixing the torsion angles 5-1-2-3 and 6-5-1-2 at 0°. 

It is found from NBO analysis that the p contribution is responsible for the nonplanar geometry and ns contribution makes the lithium bond slightly away from the p-bond center, locating it around the base atom. 

Geometry of the complexes indicates that lithium bonding is stronger in the formaldehyde complex than in the thioformaldehyde complex, and at the same time the secondary hydrogen bonding is stronger in the latter than in the former. 

NBO analysis indicates that though the donating levels vary in H2CO and H2CS complexes, electrons are accepted in the LiF antibond orbital. 

The weak interaction between closed-shell molecules plays a vital role in an enormous variety of chemical, physical, and biological phenomena. 

H2CO•••LiF has another equilibrium structure IV, and in that .C5O p-bond pair seems to be involved in the donation of charge to lithium, and here also the F•••H distance is 2.729 Å—sufficiently large to prevent the secondary hydrogen bond interaction. 

is found to be stronger than its corresponding H2CO•••HF complex reported, and this implies that lithium bonding is stronger than hydrogen bonding. 

1. There are 20 total species—10 for formaldehyde and another 10 for thioformaldehyde—the geometries of which have been optimized. 

BSSE correction at the correlated level is found to be more, and due to this, corrected binding energies at the MP2 level are found to be lower than the corresponding uncorrelated values. 

For both of the complexes, structure III is found to be the most stable, and the H2CO complex is found to be stabler by approximately 4 kcal/mol than the corresponding H2CS complex. 

NBO analysis clearly shows that this happens because sulfur donates its np lone pair while oxygen donates its ns lone pair in complexes.