From the Department of Chemistry and Biochemistry, Utah State University, Logan, Utah 84322-0300
Received for publication, November 29, 2000, and in revised form, January 3, 2001
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ABSTRACT |
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Although the peptide C Whereas conventional hydrogen bonds that involve electronegative
atoms like oxygen and nitrogen have been thoroughly studied over
the decades since their first introduction into the literature and are
presently well understood (1-3), the same cannot be said for the
CH··O interaction, which is only now gaining wide acceptance as a
genuine hydrogen bond
(H-bond)1 (4, 5). Although
normally weaker than its conventional OH··O cousin, the CH··O
interaction is thought to be crucial in a large number of molecular
complexes and crystal structures (6-10). Indeed, the CH··O bond
has been deemed so important as to foster the recommendation that the
many crystal refinement programs that treat nonbonded C··O
separations as repulsive ought to be revised (4, 11, 12).
This being the case, it would be surprising indeed if the CH··O
bond were any less important in biological systems. In fact, after some
early proposals of CH··O contacts (13-15), they were positively
identified in components like sugars (16). They are now known to be
prevalent in larger systems such as carbohydrates (17) and nucleic
acids (18-20), where these interactions can be prime determinants for
base pairing specificity (21) or general folding motifs (22). The
CH··O bond also plays an important role in the interactions of
nucleic acids with proteins (23, 24) and drug binding (25-27).
There is an increasing body of evidence that CH··O contacts occur
with some regularity in proteins as well. It was noted some time ago
that the crystal structures of various amino acids contain these
interactions (28), but their importance to larger protein segments such
as By far the most prevalent CH group in proteins involves the
C Despite the finding of numerous C It is here that one can profitably turn to ab initio quantum
calculations, a particular strength of which is the assessment of
interaction energies between various entities. In the case of the
general CH··O interaction, there have been a number of relevant
calculations (see Ref. 39 for a summary). It has been learned for
example that the interaction energy of a small prototype, the
methane-water pair, is 0.5 ± 0.1 kcal/mol (40-43). This value is
increased systematically by roughly 1 kcal/mol as each hydrogen atom of
methane is replaced by electronegative atoms like fluorine or chlorine,
which make the CH a progressively stronger proton donor (44-48). This
enhancement is not limited to simple atoms like fluorine or chlorine
but occurs as well when the CH is adjacent to larger electronegative
groups as in the context of a carboxylic acid or amide (49-53). Since
the C Reported here for the first time are the binding strengths calculated
for the C Ab initio calculations were carried out with the
GAUSSIAN 98 program using the 6-31+G** basis set (55). Electron
correlation was included by second-order Møller-Plesset perturbation
theory (MP2), which has been shown to compare favorably with more
advanced schemes for related systems (53, 56, 57). NMR chemical shift tensors were evaluated by the gauge-independent atomic orbital (58) method. The binding energies are computed as the difference in
total energy between the complex on one hand and the sum of the
isolated, optimized monomers on the other; basis set superposition error is removed by the counterpoise procedure (59).
Gly, Ala, and Val are taken as representative of the nonpolar amino
acids. Ser and Cys contain the polar OH and SH groups, respectively. As
examples of charged residues, the Lys+ cation and the
Asp Binding Energies--
The interaction energies of each of
the various amino acids with water as the proton acceptor are reported
as
The replacement of the two fluorine atoms by the COOH and
NH2 groups, respectively, results in the
NH2CH2COOH amino acid glycine. Since the two
substituent groups are rather electronegative, much as the two fluorine
atoms, one might anticipate only a minor perturbation upon the binding
energy in the complex with water. In fact, inspection of Table I
confirms this expectation in that glycine and F2HCH have
nearly identical H-bond energies. The remainder of Table I focuses upon
the changes conferred by replacement of glycine by each of several
other amino acids. Substitution of one hydrogen atom by methyl yields
the alanine residue and a reduction of the interaction energy by 0.4 kcal/mol. Enlargement of the methyl group of alanine to the isopropyl
group of valine reduces the binding energy by another 0.1 kcal/mol. It
is logical to presume that the slightly larger aliphatic side chains of
Leu and Ile would be similar to the result for Val, and that the
binding energies of this series of amino acids, containing simple alkyl
side chains, lie in the range between 2.0 and 2.5 kcal/mol.
The serine residue contains the polar hydroxyl group in its
CH2OH side chain. Nonetheless, its calculated H-bond energy
of 2.3 kcal/mol falls right within the range of the nonpolar residues. In a more highly refined way of looking at the data, the replacement of
one hydrogen atom of the Ala methyl group by the more electronegative OH enhances the H-bond energy by some 0.2 kcal/mol. In contrast, the
CH2SH sulfhydryl side chain of the Cys residue reduces the binding energy of Ala by 0.2 kcal/mol. In summary, the H-bond energies
of the above amino acids, including aliphatic side chains, polar
CH2OH, and less polar CH2SH, are all quite
similar to one another, in the 1.9-2.5 kcal/mol range.
Moving on to the charged residues, we first consider the cationic
lysine residue with its
(CH2)4NH
Turning now to aspartate, the upper curve of Fig.
1 illustrates the potential for the
interaction between the aspartate residue (R=CH2COO Other Aspects of Interaction--
The second column of Table I
lists the equilibrium distances between the proton donor atom and the
oxygen of the water acceptor, the intrinsically preferred H-bond
length. This quantity is generally correlated with
The effect of the formation of the CH··O H-bond upon the CH bond
length is reported in the third column of Table I (in mÅ). The first
row illustrates the 5-mÅ stretch in the water dimer, a stretch that is
characteristic of conventional OH··O bonds. This elongation
behavior contrasts markedly with the contractions that occur in the
CH··O H-bonds of all amino acids as well as F2HCH. This
seemingly opposite behavior, observed in a number of CH··O bonds
(46, 62-67), has nevertheless been demonstrated to be consistent with
the characterization of the CH··O interaction as a true H-bond
(54). There is no clear pattern concerning the magnitude of this
contraction. For example, Gly and Val elicit very small contractions,
while a much larger one occurs in the similar Ala. The positive charge
of the Lys+ residue does not appear to unduly raise the
magnitude of this bond shortening. The largest contraction of all is
associated with the aspartate anion, although its complex with water
might be considered the weakest of all those considered. At any rate, the consistency of the negative values in all cases supports the idea
that the C
Associated with the OH bond stretch of a conventional H-bond is a red
shift of its vibrational frequency. As listed in the first row of Table
I, this shift amounts to
As mentioned above, these opposing patterns in the OH and CH covalent
bond behavior are not entirely unexpected, having been observed on a
number of occasions. There is now some reason to believe that the CH
bond contraction/blue shift commonly occurs when the C atom is bonded
to four other atoms, i.e. sp3 hybridization, as
in the amino acids, but that the sp hybridization of the alkyne CH is
associated with the stretch and red shift typical of conventional
H-bonds (39). The magnitude of blue shifts calculated here is
consistent with prior work dealing with CH··O bonds (67, 75). In
any event, these blue shifts ought to serve as a marker of the presence
of a C
The next column of Table I reports the effect of H-bond formation upon
the calculated intensity of each CH bond stretching frequency. The
water dimer undergoes a characteristic enhancement, with the band 1.9 times stronger in the complex than in the isolated water monomer. There
is little obvious pattern within the CH··O complexes, in that some
bands are intensified (value greater than 1) while others behave in the
opposite fashion. Hence, while a blue shift in frequency can be taken
as a clear indication of the formation of a CH··O bond, it might be
misleading to use the intensity as an indicator.
Nuclear magnetic resonance frequencies have been used to good effect to
monitor the presence of hydrogen bonds (76). The calculated values of
the shifts of the bridging hydrogen are reported in the last two
columns of Table I, relative to the isolated monomers. The isotropic
shift is listed first, followed by the anisotropic value. Probably the
most well recognized NMR diagnostic of the presence of a H-bond is a
downfield (negative) isotropic shift of the bridging hydrogen (77).
This quantity is calculated here to be The results presented here were obtained at the MP2/6-31+G**
level, which, as pointed out above, there is reason to believe is quite
reliable for such systems. For example, our own earlier work on related
CH··O systems (54) had demonstrated a rather remarkable
insensitivity to details of the basis set. A similar test was carried
out here, wherein a set of diffuse functions was added to the hydrogen
atoms, so as to better match the basis sets on the heavier atoms. The
MP2 binding energy computed with this larger 6-31++G** basis set for
the Gly-water complex matched within 0.02 kcal/mol the result listed in
Table I obtained with the smaller set. Other aspects of the
interaction, viz. the optimized H-bond length, the change in
the CH bond length, and its associated blue shift, were also virtually
unchanged by this basis set enlargement.
The water molecule has been taken as the model proton acceptor in the
hydrogen bonds discussed here. While HOH is in fact one of the acceptor
molecules that one would expect to participate in such interactions, it
also adequately mimics the hydroxyl group that occurs on such residues
as serine or threonine. However, the situation also arises where the
oxygen acceptor is involved in a double bond as in the peptide group or
asparagine. It is thus important to consider how the change in bonding
pattern from hydroxyl -OH to C=O affects the strength of the H-bond.
Earlier calculations in which CFnH4-n had
been used as proton donor (54) had demonstrated that the binding to an oxygen acceptor of the C=O sort is only slightly weaker than the H-bond
to water. The same was found here, with amino acids acting as the
donor. Taking our glycine residue as proton donor, the replacement of
the water acceptor by H2CO reduces the binding energy from
2.5 to 2.3 kcal/mol. The donor's 1-mÅ CH bond contraction is observed
with H2CO as well, as is the blue shift in its stretching frequency. Nor does the change of acceptor from HOH to H2CO
have much effect upon the isotropic NMR shift of the bridging hydrogen, although there is a significant reduction in the anisotropic shift.
Whereas the C Some earlier calculations on related systems present some basis for
comparison, although the complexes studied contain several distorted
CH··O H-bonds, complicating the estimation of the H-bond energy of
any single undistorted H-bond. In acetic acid, the proton-donating methyl group abuts a carboxyl, as opposed to an amide. A study of the
acetic acid dimer (49) arrived at an estimate of each CH··O H-bond
energy in the range of 0.5-1.8 kcal/mol, which may be considered a
lower bound due to the aforementioned distortions. Later calculations
obtained CH··O bond energies of the aldehydic CH with various
proton acceptors between 0.8 and 2.6 kcal/mol (51). A calculation of
the formamide dimer (50) was somewhat more relevant in that the CH
group is part of an amide; the bond strength was placed in the 2.5-4.0
kcal/mol range. While Vargas et al. (52) did not study any
amino acids or protein residues directly, they did compute an accurate
value for the total interaction energy of a complex involving a pair of
ordinary amides, for which the binding CH was part of a
terminal-CH3 group. As in the other work, these complexes
each contained a variable number of distorted CH··O H-bonds, some
obviously stronger than others, so the H-bond energy of any single
undistorted H-bond can only be guessed. The determination of the energy
of a single H-bond was further complicated by cooperativity effects
that act to strengthen the interaction. These problems notwithstanding,
the authors arrived at H-bond energies in the range between 2.1 and 2.7 kcal/mol, only slightly larger than the values computed here for true
amino acids, further support for the accuracy of our data.
It is emphasized that the binding energies reported above refer not to
free energies or enthalpies but rather to the electronic contribution
to the interaction energy. One can convert the latter quantity to
enthalpy at a physiologically relevant temperature by incorporating
vibrational energies, as well as other corrections. Doing so provides
values of As mentioned earlier, the structures of the various complexes have been
optimized under the restriction of a linear CH··O arrangement. As a
result, the optimized complex is not, strictly speaking, a true minimum
on the entire potential energy surface. As expected, relaxation of this
restriction permits the water molecule to swing around toward the COOH
group, forming an H-bond between the carbonyl oxygen of the COOH and
one of the water hydrogens, a bond that is stronger than the CH··O
interaction of interest. Nonetheless, even in the presence of this
stronger OH··O interaction, the properties of the
C Concerning the relative strengths of the two sorts of H-bonds, it may
be worth noting that an earlier set of calculations had demonstrated
that CH··O H-bonds are less sensitive to geometrical distortions
than are conventional OH··O interactions (54); i.e. the
strength of a CH··O interaction is eroded more gradually by stretching or bending from its optimal configuration than are OH··O
bonds. Hence, even if the former is intrinsically weaker than the latter (when both are in their preferred geometry), the situation can reverse with the stretches or bends, which are the rule
rather than the exception in proteins. Taking into account the
appreciable intrinsic strength of the CH··O bond and its resistance to weakening via bond distortion, this interaction cannot be discounted when considering the various factors that lead to protein structure. Indeed, the occurrence of such bonds may be an important factor in the
presence of water molecules within the confines of proteins.
The question of whether the C In summary, the CH
group has historically not been thought to form hydrogen bonds within
proteins, ab initio quantum calculations show it to be a
potent proton donor. Its binding energy to a water molecule lies in the
range between 1.9 and 2.5 kcal/mol for nonpolar and polar amino acids;
the hydrogen bond (H-bond) involving the charged lysine residue is even
stronger than a conventional OH··O interaction. The preferred
H-bond lengths are quite uniform, about 3.32 Å. Formation of each
interaction results in a downfield shift of the bridging hydrogen's
chemical shift and a blue shift in the C
H stretching
frequency, potential diagnostics of the presence of such an H-bond
within a protein.
INTRODUCTION
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ABSTRACT
INTRODUCTION
THEORY
RESULTS
DISCUSSION
REFERENCES
-glycine (29) has been revealed as well. Other groups that appear
to be involved in CH··O H-bonding include the aryl groups of
aromatic residues like phenylalanine (30), the C
H of
proline (31), the CH groups of histidine (32, 33), and the lysine
C
H and valine C
2H groups
(34).
of each amino acid residue, so its possible
involvement in H-bonds is of profound consequence. Even if individually
weak, the sheer number of such C
H··O H-bonds could
exert an enormous influence upon the structure and function of a
protein (4, 35). Perhaps the earliest direct evidence that the
C
H group might in fact participate in H-bonds derives
from a neutron diffraction study of amino acid crystals (28), which
found geometric indicators of as many as 16 different
C
H··O H-bonds. This idea was later confirmed in a
wide variety of proteins (36), including the collagen triple helix
(37), and in
-sheets (38), where the CH··O bonds are thought to
confer additional stability.
H··O contacts in
proteins, major questions remain about their importance. While
providing a wealth of information about geometries, the numerous
crystal structures on which most current knowledge about the
C
H··O interaction is based are silent on the
energetic aspects. In short, "nothing is known experimentally about
the strength of these interactions" (35). Yet it is the strength of
this binding that is of most importance in understanding the possible role that the C
H··O H-bond may play in the folding
and function of proteins.
H group of each amino acid residue in a protein is
directly adjacent to a pair of electronegative groups (the N and C ends
of two amide groups), it is logical to presume that its ability to form
an H-bond is comparable with that of a small molecule like
CH2F2. Since recent calculations (54) have
demonstrated that the latter molecule does form a true H-bond, which,
under certain circumstances, can be of a strength similar to a
conventional OH··O interaction, an explicit examination of the
H-bonding abilities of the C
H group of real amino acid
residues is warranted.
H group of a number of representative amino
acids, with a common oxygen acceptor. The results indicate the
comparative H-bond energy of each. In addition, supplementary IR and
NMR spectroscopic information are computed so as to aid in the
identification of such bonds in an experimental setting.
THEORY
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ABSTRACT
INTRODUCTION
THEORY
RESULTS
DISCUSSION
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anion were considered. All amino acids were
considered in their NH2CHRCOOH nonzwitterion state so as to
better model their neutral condition within a protein. Water was taken
as the proton acceptor in each complex. Geometries were fully
optimized; the sole restriction was that a
(CH··O) angle of
180° was maintained in the complex to prevent the formation of
complicating secondary interactions.
RESULTS
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ABSTRACT
INTRODUCTION
THEORY
RESULTS
DISCUSSION
REFERENCES
E in the first column of data in Table
I, under the convention that a negative
E corresponds to a favorable binding energy. The first
row illustrates data for the water dimer, as a classic paradigm of an
OH··O hydrogen bond, for which the electronic contribution to the
binding energy is 4.51 kcal/mol at this level of theory. (This value
compares quite favorably with quantities predicted at higher levels
(60).) Replacement of the OH of the water donor by a CH group is
expected to weaken the H-bond considerably. This supposition is correct in that even after the addition of two electronegative fluorine atoms
to enhance its acidity, proton donor F2HCH binds to the water acceptor by only 2.53 kcal/mol, about half the value for the
water dimer (54), as reported in the next row of Table I.
MP2/6-31+G** calculated properties of interaction of various proton
donors with water
H-O interaction.
H group a more powerful proton donor, enhancing its
H-bond to water, as occurs with conventional H-bonds. This expectation
is verified in the case of the Lys+ residue, with a binding
energy of 4.9 kcal/mol, about double that of the neutral residues,
although the -NH3+ group bearing the formal
charge is several bonds removed from the site of H-bonding. In fact,
this particular CH··O H-bond is slightly stronger than the
conventional OH··O of the water dimer.
) and a water molecule. It may be
first noted that this curve contains a minimum at a R(C··O)
separation of about 3.33 Å. A stretch of this H-bond by as much as 2 or 3 Å acts against an attractive force pulling the two groups
together. However, once the separation has reached 5 or 6 Å and the
energy has risen to 2.5 kcal/mol above the minimum, the force changes
from attractive to repulsive, acting to push the groups even further
apart. This long range repulsion is understandable on the basis of the
electrostatic repulsion between the aspartate anion and the negative
end of the water molecule's dipole moment. In one respect, this
interaction represents a H-bond of about 2.5 kcal/mol, since that is
the energy required to overcome the barrier in the potential. (One must
be cautious about the definition of the H-bond energy in such a case, since Table I reveals that the complex is higher in energy than the two
separated monomers by 1.4 kcal/mol.) The situation is less complex in
the cases of the other amino acids where the potentials have no
maximum. The behavior of the complex between Ala and water, illustrated
by the lower curve in Fig. 1, is a case in point, where the H-bond energy is defined simply as the energy of the minimum,
relative to infinite separation. In conclusion, there appears to be an
attractive interaction that prevents the separation of the aspartate
residue from water, despite their long range repulsive interaction. It
is intriguing, but undoubtedly coincidental, that the height of this
barrier is very close in magnitude to the H-bond energies of the
neutral amino acids.
View larger version (11K):
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Fig. 1.
Interaction energies as a function of
intermolecular separation between water and amino acids alanine and
aspartate. Energies have not been corrected by the counterpoise
procedure.
E,
with stronger H-bonds associated with a shorter length. It is therefore
interesting to find that this H-bond distance is rather uniform in all
CH··O H-bonds, covering a rather narrow range between 3.31 and 3.35 Å. This range agrees quite well with the H-bond length of 3.35 Å measured by neutron diffraction for the interaction between a
C
H group and a water molecule (61), as well as the
average C
··O distance of 3.31 Å in parallel
-sheets (38). The H-bond length is nearly constant, about 3.34 Å,
for F2HCH as well as for the three aliphatic amino acids
Gly, Ala, and Val. It shortens slightly to 3.31 Å for the Ser and Cys
residues. Despite the stronger binding of the Lys+ residue,
its H-bond length of 3.32 Å is in the same range as the other
complexes, as is the separation (3.33 Å) for the anionic Asp
.
H bond is shortened, albeit by a small amount,
in all amino acid H-bonds, as occurs in a variety of other CH··O interactions.
31 cm
1 for the
water dimer. The positive signs of the remaining entries in the fourth
column of Table I indicate that the amino acid CH bonds all shift in
the opposite direction, to the blue. These positive shifts correlate
with the CH bond contractions and are in fact consistent with a number
of experimental observations involving CH··O bonds over the years
(68-74). Also consonant with previous findings, the magnitudes of
these shifts are variable and, like the contractions of the CH bond, do
not fit a simple pattern, although there is a trend for larger blue
shifts to be associated with a greater amount of bond contraction. The
largest shift of 70 cm
1 is associated with
the aspartate anion.
H··O H-bond in proteins.
2.6 ppm for the water dimer
and lies in the range between
1.3 and
1.7 for the various amino
acids. This range is consistent with computations of other CH··O
interactions (78, 79) as well as experimental measurements in solvent
(69, 80-82) or protein environment (33). The anisotropic shifts of the
hydrogen (last column) are also rather uniform, occupying the 6-8-ppm
range for the CH··O bonds, in comparison with 11 for the OH··O bond.
DISCUSSION
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ABSTRACT
INTRODUCTION
THEORY
RESULTS
DISCUSSION
REFERENCES
of an amino acid is surrounded by
NH2 and COOH groups, it lies adjacent to full peptide
groups within the context of a protein. To gauge how the results might
be affected when the residue is part of a protein, the NH2
and COOH groups of the glycine amino acid were both enlarged to full
amide groups, creating the larger
HCONHCH2CONH2. The binding energy of this model
with water was greater than that of the Gly amino acid by 0.3 kcal/mol. One can assume therefore that the binding energies for each of the
amino acids in Table I would be increased by a like amount when the
residue is surrounded by peptide groups. The CH··O H-bond energies
ought therefore to fall into the range 2.5 ± 0.3 kcal/mol for the
uncharged amino acids, both polar and nonpolar, but as much as 5.2 kcal/mol for a cationic residue such as Lys+.
H298 that are small enhancements of
the binding energies in Table I, varying between
3.0 and
3.8
kcal/mol for Gly through Cys. Whether
E or
H, it is important to stress that our computed values
refer to an intrinsic binding energy of the amino acid C
H··O H-bond for purposes of comparison with the
conventional OH··O interaction; external effects of the protein
environment or of solvation have not been explicitly included for
either sort of H-bond but are expected to be similar for both.
H··O H-bond remain qualitatively the same. Taking
the complex between Gly and water as an example, the full optimization
bends the C
H··O bond from 180 to 136°, but the
C
H bond is shortened by the interaction in either case.
The amount of this contraction is 1.4 mÅ for the fully optimized
minimum, as compared with 1.0 mÅ when C
H··O is held
linear; the corresponding blue shifts of the C
H stretch
are 9 and 14 cm
1, respectively.
H··O interaction
represents a genuine H-bond is not restricted to the binding energy but
involves other aspects of the phenomenon as well. Earlier work (54)
made use of ab initio calculations to document the
similarities in electron density shifts that occur upon formation of
the CH··O and OH··O interactions. It was noted there, for
example, that the charge transfer from one subunit to the next obeys
similar patterns in the two sorts of interactions, as do the dipole
moment enhancements of the complex, compared with the isolated
monomers. Also quite similar are the changes in the charges of the
various atoms caused by the interaction, as well as contour maps that illustrate charge shifts over all regions of space. Another feature that the two sorts of H-bonds share is the contribution of electron correlation to the binding. In the water dimer, for example,
correlation contributes as much as 0.6 kcal/mol, depending upon the
particular level of theory. Rather similar contributions were obtained
here for the interactions between the various amino acids and water, in
the range between 0.3 and 0.7 kcal/mol; it should be noted that the
latter represent larger percentage contributions in the C
H··O case, due to the generally weaker binding as
compared with OH··O.
H··O interaction appears to be a
true H-bond. The binding energy for the equilibrium separation is in
the neighborhood of 1.9-2.5 kcal/mol for the uncharged residues,
roughly half that of the water dimer; the H-bond energy of the
Lys+··OH2 pair is even stronger than the
conventional OH··O bond. These observations suggest that the
C
H··O bond must be considered a potentially
important factor in protein structure and function. The optimum
R(C··O) separation, for maximum binding strength, is some
3.31-3.35 Å. All of these H-bonds undergo a contraction of the CH
bond, with a characteristic blue shift of its stretching frequency that
may serve to help identify their presence. The NMR shifts of the
bridging proton, another useful diagnostic, are characteristic of
conventional H-bonds.
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FOOTNOTES |
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* This work was supported by National Institutes of Health Grant GM57936.The costs of publication of this article were defrayed in part by the payment of page charges. The article must therefore be hereby marked "advertisement" in accordance with 18 U.S.C. Section 1734 solely to indicate this fact.
To whom correspondence should be addressed. Tel.: 435-797-7419;
Fax: 435-797-3390; E-mail: scheiner@cc.usu.edu.
§ Present address: 175 E. Park Dr., Praxair, Tonawanda, NY 14151.
Published, JBC Papers in Press, January 4, 2001, DOI 10.1074/jbc.M010770200
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ABBREVIATIONS |
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The abbreviation used is: H-bond and H-bonding, hydrogen bond and hydrogen bonding, respectively.
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REFERENCES |
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