Hydrogen, bonding

Hydrogen bonds are the most characteristic element of liquid water structure. Water models used in computer simulations are able to describe the properties of the hydrogen bond network in a realistic way, contrary to many of the dipolar model fluids used in analytical theories. Much has been learned about bulk water and solutions through an analysis of the hydrogen bond network (e.g.. Ref. 156, 157). 

In Fig. 20 the local number of hydrogen bonds per molecule is plotted versus the local coordination number at all distances and for several simulations. There is a clear correlation between the decrease of the number of hydrogen bonds and the increase in coordination number for all simulations, independent of the interface. The values near the interface (at large nn) fall on the same line as the bulk values. 

The hydrogen bond (A-H B) is mainly electrostatic in character. Its strength appears to be a compHcated balance of various factors [28] which include  

the interactions between fractional charges that have developed on A, B, and H, 

the deformabihty (polarizability) of the electron cloud aroimd the acceptor atom B so that it can make its lone pairs available to the proton (the softness ofB), 

the transfer of electronic charge from B to H (n-bond transfer), 

how readily a hydrogen-bond donor atom A will lose its covalently boimd hydrogen atom as H+ (related to the electronegativity of A and the strength 

In a hydrogen bond, here written H-bond, a hydrogen atom occupies a more or less central position between, usually, two electronegative atoms. This is conventionally represented as a linear A-H—B combination of the hydrogen donor system, A-H, and the acceptor, B. Relevant properties of the different strengths of H-bonds are given in Table 9.1 [1]. 

Properties Strong H-bonds Moderate H-bonds Weak H-bonds 

Bond nature mostly covalent mostly electrostatic electrostatic 

Bond linearity, A-H— -B always linear mostly linear sometimes linear 

This chapter focuses on results from INS for the stronger and moderate H-bonds where, as explained below, the dynamics of hydrogen transfer should be accessible. Since most known examples of moderate H-bonds occur in oxygen-oxygen systems they will assume a particular prominence in this chapter. H-bonds also determine, to a great extent, the tertiary and quaternary structure essential to achieve biological activity ( 10.3). 

These hydrogen bonds are of lower strength than the covalent bonds in the rest of the chains but are still sufficient to form a strong compound. The hydrogen bond is approximately 5 kcals/mol, compared to the covalent bond of about 50 to 100 kcals/mol. 

Bulky side chains in the hard segment will tend to make the hydrogen bonding more difficult. The bulky side chains of Ethacure 300 will lower the hardness of a material by several points compared to that of a material cured with MOCA. The use of triols or macro diols in the curing phase will have a softening effect. If a trifuctional isocyanate or hydroxyl is used in the initial preparation of the prepolymers, this softening does not happen. 

Using a straight polyurethane system, for example, a PTMEG/MDI prepolymer, there will initially be an absorption band at approximately 1733 to 1725 cm1 (Seymour et al., 1970). As the cure starts, this band will initially 

As a result of the shape of the molecules and the fact that they are twisted and bent, the bonding is in all three planes and not as a flat object. 

Different references give variations to the values for the different bands. These vary depending on the calibration of the instrument used, the exact chemistry of the system, and the method of preparation and testing of the sample. Samples tested in a solvent may indicate the carbonyl band at a wave number of 1745 cm-1, whereas the solid material will be at approximately 1730 cm-1. 

In the hydrogen chloride molecule there is a permanent dipole. Draw this molecule showing the direction of this dipole. 

The lone pairs of the chlorine atom are electron rich and so may become involved in an electrostatic interaction with the 8 positive centre of the polarised hydrogen atom. 

This Coulombic interaction is called a hydrogen bond, and is a source of stabilisation in a variety of molecules, in particular in those that have an atom bearing a negative charge that is close to a polarised hydrogen atom. For example, such a configuration is present in ortho-hydroxybenzoic acid after the carboxylic acid group has been deprotonated. Draw the resultant anion. 

In this case there is an intramolecular hydrogen bond, while in the case of the hydrogen chlorine molecule the hydrogen bond was intermolecular. In the case of o/v/ o-hydroxybenzoic acid, hydrogen bonding is also present in the neutral acid, but to a lesser extent. 

In dicarboxylic acids that are capable of forming intramolecular hydrogen bonds, it is observed that the first proton is removed more easily than would be predicted by just considering the inductive effect of the second carboxylic acid. Suggest the structure that c/.s-butenedioic acid adopts after one deprotonation. 

Hydrogen bonding can occur in any system containing a proton donor group (X—H) and a proton acceptor 

FIGURE 3.3. Infrared spectrum of cyclopentanone in various media. A. Carbon tetrachloride solution (0.15 M). B. Carbon disulfide solution (0.023 M). C. Chloroform solution (0.025 M). D. Liquid state (thin films). (Computed spectral slit width 2 cm-1.) 

The strength of the hydrogen bond is at a maximum when the proton donor group and the axis of the lone pair orbital are collinear. The strength of the bond decreases as the distance between X and Y increases. 

X—H Y Strength Frequency Reduction (cm-1) OH C=0 Compound Class Frequency Reduction (cm-1) VOn VC=0 Compound Class 

Weak Medium 300 1 15 Alcohols, phenols, and intermolecular hydroxyl to carbonyl bonding 100° 10 100-300 50 1.2- Diols, a- and most /3-hydroxy ketones o-chloro and oal-koxy phenols 1.3- Diols some /3-hydroxy ketones /3-hydroxy amino compounds /3-hydroxy nitro compounds 

FIGURE 2.3. Infrared spectrum of cyclopentanone in various media. A. Carbon tetrachloride solution (0.15 M). 

Weak 300 15ft Alcohols, phenols, and inter- 100 10 1,2-Diols, a- and most /3-hydroxy 

Strong 500 50fr RC02H dimers 300 100 o-Hydroxy aryl ketones o-hy- 

Hydrogen bonding is important throughout chemistry and biochemistry. Many gas-phase hydrogen-bonded dimers have been characterized spectroscopically examples include (H20)2, (HC1)2, HF-H2O, and HF-HCN. For (H20)2, the structure (determined by molecular-beam microwave spectroscopy) is [T. R. Dyke, K. M. Mack, and J. S. Muenter,/ Chem. Phys., 66,498 (1977)] shown in Fig. 17.1. 

For H-bonded dimers, HF STO-3G, 3-21G, and 3-21G calculations give equilibrium-geometry separations between the heavy atoms that are usually substantially in error (errors of 0.1 to 0.5 A) 6-31G calculations give heavy-atom separations in pretty good agreement with experiment Hehre et al.. Table 6.32). 

However, this procedure involves an inconsistency. When the monomer energy A( fA ) is calculated, the electrons of A have available to themselves only the 3-21G orbitals on the three atoms of A, whereas, when eAB( fA + fB ) is calculated, the electrons of each H2O molecule within the dimer have available not only the orbitals on their own nuclei, but also the orbitals on the nuclei of the other H2O molecule. In effect, the dimer basis set is larger than that of each monomer, and this produces an artificial lowering of the dimer energy relative to that of the separated monomers. This artificial lowering is called the basis-set superposition error (BSSE). The BSSE would vanish in the limit of using a complete set for each monomer. The most often used procedure to correct for the BSSE is to calculate the dimerization energy as 

From the temperature and pressure dependences of the thermal conductivity of water vapor. A//373 for 2H20(g) (H20)2(g) has been found to be -3.6 0.5 

HF STO-3G, 3-21G, 6-31G, and 6-31G calculations give H2O dimerization electronic energies of -5.9, -11.0, -5.6, and -5.5 kcal/mol, respectively, and counterpoise-corrected dimerization energies of -0.2, -6.2, -4.65, and -4.5s kcal/mol, respectively [M. J. Frisch et al., /. Chem. Phys., 84, 2279 (1986)]. SCF calculations with huge basis sets and the CP correction give the dimerization energy in the Hartree-Fock limit as -3.73+0.05 kcal/mol. 

To form a hydrogen bond what must the non-hydrogen atom (N, O, or F) involved in the bond possess  

Hydrogen bonds can be considered a type of dipole-dipole attraction. Because N, O, and F are so electronegative, a bond between hydrogen and any of these elements is quite polar, with hydrogen at the positive end (remember the -F on the right-hand side of the dipole symbol represents the positive end of the dipole)  

Identifying Substances That Can Form Hydrogen Bonds 

In which ot these substances is hydrogen bonding likely to play an important role in determining physical properties methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), hydrogen sulfide (H2S)  

Analyze We are given the chemical formulas of four compounds and asked to predict whether they can participate in hydrogen bonding. All the compounds contain H, but hydrogen bonding usually occurs only when the hydrogen is covalently bonded to N, O, or F. 

When hydrogen bonding occurs between two molecules, the molecules both have to be polar. Therefore, in addition to hydrogen bonding, dipole-dipole forces are also present. 

Nonpolar molecules do not contain a dipole, at least a fixed dipole. However, there are forces of attraction between nonpolar molecules. Otherwise, nonpolar gases like CO, N, O, and He would never condense into liquids. The fact that at sufficiently cold temperatures a nonpolar gas does condense means there must be some kind of intermolecular force of attraction between its molecules. That force is too weak at high temperatures for the gas to condense, but at cold temperatures, it is strong enough to cause condensation. The same force is responsible for liquids freezing into solids at even colder temperatures. 

This weak intermolecular force is caused by fluctuations in the electron distribution in a molecule that creates temporary, or instantaneous, dipoles. These instantaneous dipoles appear and disappear on a very short time scale. On a longer time scale, there is no fixed dipole and the molecule is nonpolar. However, these instantaneous dipoles exert just enough force on surrounding molecules to hold the molecules in the liquid or solid phase. This force 

The electrical charges in permanent dipoles fluctuate slightly in their distribution. Vhen dipole-dipole forces are present between two molecules, those fluctuations also cause London forces to be present. In general the London forces will be weaker than the dipole-dipole forces, but both contribute to the attractions between molecules. If a molecule also exhibits hydrogen bonding, dipole-dipole and London forces will both be present and contribute to the attractions between molecules. 

The discovery of the hydrogen bond was not a historical event Uke, for example, the discovery of x-rays. Rather, it was gradually realized that hydrogen binds weakly to electronegative atoms other than those it is bonded with in normal chemical bonds (Latimer and Rodebush 1920). Often, quite strong intermolecular bonds are formed. The PT in bond has become a very important subject, since the processes of life rely on coupled electron and proton transfer. 

FIGURE 9.2 Hydration around the ammonium ion NHJ. The water molecules of the first solvent shell are connected to the nitrogen atom by dashed lines. 

Because intermolecular hydrogen bonds between small molecules are usually made and broken very rapidly relative to the chemical shift difference between bonded and nonbonded forms (expressed in hertz), separate lines are generally not observed for different species, and the frequency v of the single observed line is given by Eq. 2.58. In suitable systems the variation of v with concentration and temperature can be analyzed to give equilibrium constants and other thermodynamic properties for hydrogen bond formation. 

FIGURE 4.12 H NMR spectrum (500 MHz) of a self-complementary decadeoxynucleotide, (/(GCATTAATGC), in H20 at a concentration of 20 millimolar. From Wiithrich.60 

NMR has become an indispensable analytical technique for studying rates of H D exchange at specific sites when a protein is deuterated. Such exchange rates provide information on secondary protein structure and are useful in investigating the mechanism of protein folding. 

The Chemical Bond Chemical Bonding Across the Periodic Table, First Edition. 

Student Annotation Flydrogen bonding also occurs in mixtures, between solute and solvent molecules that contain N—FI, F—FI, or 0—FI bonds. 

CHAPTER 12 Intermolecular Forces and the Physieal Properties of Liquids and Solids 

The eoneept of hydrogen bonding was introdueed in 1919 by Huggins [37], The first definitive paper on hydrogen bonding - applied to the assoeiation of water moleeules - was published in 1920 by Latimer and Rodebush [191], All three were working in the Laboratory of G. N. Lewis, University of California, Berkeley/USA. 

A hydrogen bond is formed by the interaetion between the partners R—X—H and Y—R aeeording to Eq. (2-7). 

When two or more molecules of the same type associate, so-called homo-intermolecular hydrogen bonds are formed (Fig. 2-4). The association of different molecules e.g. R—O—H- -NRj) results in hetero-intermolecular hydrogen bonds. The designations homo- and heteromolecular [192] as well as homo- and heteroconjugated hydrogen bond are also in use. A remarkable example of a competitive solvent-dependent equilibrium between homo- and hetero-intermolecular hydrogen-bond associated species has been found in solutions of 4-hydroxyacetophenone and 2-(2-hexyloxyethoxy)ethanol [319]. 

Hydrogen bonds can be either intermolecular or intramolecular. Both types of hydrogen bonds are found in solutions of 2-nitrophenol, depending on the Lewis basicity of the solvent [298]. The intramolecularly hydrogen-bonded form exists in nonhydrogen-bonding solvents e.g. cyclohexane, tetrachloromethane). 2-Nitrophenol breaks its intramolecular hydrogen bond to form an intermolecular one in electron-pair donor (EPD) solvents e.g. anisole, HMPT). 

The question of the exact geometry of hydrogen bonds (distances, angles, lone-pair directionality) has been reviewed [194]. 

JJ Would you expect chloroform to form hydrogen bonds If so, would those bonds be intramolecular or intermolecular  

WouM yqu expecyfie CP-H stretchmg nigde to afy jtnfensi J en hydrogen If so  

Many solvents ar,e capable of forming hydrogen bonds to solutes. In order to form a hydrogen bond we need a proton donor group and an electron donor group. Three classes of solvent exist which may cause problems when used for hydrogen-bonding studies in solution  

Compounds which contain hydrogen donor groups, e.g. halogenated compounds which contain a sufficient number of halogens to activate the hydrogens present, such as chloroform. 

Compounds which contain non-bonded electron pairs, such as ethers, aldehydes, and tertiary amines. 

In some cases, hydrogen bonding can occur between one molecule in which H is directly bonded to F, O, or N and another molecule containing an electronegative atom. (See the box, Chemistry and Health, in this section for an example.) 

One of these compounds is a liquid at room temperature. Which one and why O 

The three compounds have similar molar masses. 

CAN YOU ANSWER THIS INhy would dispersion forces not work as a way to hold the two halves of DNA together Why would covalent bonds not work  

A FIGURE 12.25 The structure of DNA DNA is composed of repeating units called nucleotides. Each nucleotide is composed of a sugar, a phosphate, and a base. 

When two compounds whose molecules form hydrogen bonds with each other are both dissolved in water, the hydrogen bond between the two molecules is usually greatly weakened or completely removed,9 because the molecules generally form hydrogen bonds with the water molecules rather than with each other, especially since the water molecules are present in such great numbers. 

7For reviews, sec Abraham Doherty Kamlet Taft Chem. Br. 1986, 551-554 Kamlet Abboud Taft Prog. Phys. Org. Chem. 1981,13, 485-630. For a comprehensive table and a and p values, see Kamlet Abboud Abraham Taft J. Org. Chem. 1983, 48, 2877. For a criticism of the p scale, see Laurence Nicolet Helbert J. Chem. Soc., Perkin Trans. 2 1986, 1081. See also Nicolet Laurence Lugon J. Chem. Soc., Perkin Trans. 2 1987, 483 Abboud Roussel Gcntric Sraidi Lauransan Guihgncuf Kamlet Taft./. Org. Chem. 1988,53, 1545 Abraham Grcllier Prior Morris Taylor J. Chem. Soc., Perkin Trans. 2 1990, 521. 

12A statisical analysis of x-ray crystallographic data has shown that most hydrogen bonds in crystals are nonlinear by about 10 to 15° Kroon Kanters van Duijneveldt-van de Rijdt van Duijneveldt Vliegenthart J. Mol Struct. 1975, 24, 109. See also Ceccarelli Jeffrey Taylor J,- Mol Struct. 1981, 70, 255 Taylor Kennard Versichel J. Am. Chem Soc. 1983, 105, 5761 1984,106, 244. 

In certain cases x-ray crystallography has shown that a single H—A can form simultaneous hydrogen bonds with two B atoms (bifurcated or three-center hydrogen bonds). An example is an adduct (1) formed from pentane-2,4-dione (in its enol form) and diethylamine, in 

ITHine Ahn Gallucci Linden J. Am. Chem. Soc. 1984, 106, 7980 Hine Hahn Miles J. Org. Chem. 1986, 51, 

When H bonds directly to F, O, or N, the bonding atoms acquire relatively large partial charges, giving rise to strong dipole-dipole attractions between neighboring molecules. 

In HF The hydrogen of one HF molecule, with its partial positive charge, is attracted to the fluorine of its neighbor, with its partial negative charge. This dipole-dipole interaction is an example of a hydrogen bond. 

The partially positive charge on H is strongly attracted to the partial negative charge on 0. 

A FIGURE 11.11 Hydrogen Bonding in Ethanoi The left side shows the space-filling models, and the right side shows the electrostatic potential maps. 

In some force fields hydrogen bonding is handled simply by the electrostatic term just introduced. In others, there is an explicit equation for hydrogen bonds. One form for such an equation is Eq. 2.49. 

This equation is a derivative of a Lennard-Jones potential function. In this equation Nh is the number of hydrogen bonds, while C and D are parameters depending on the type of hydrogen bond. In this approach we have to explicitly define all the hydrogen bonds in advance so this equation can be applied to them. When the simple electrostatic approach is used, hydrogen bonds need not be defined explicitly. 

In order to formulate a valid constraint reflecting hydrogen bond plausibility, the first adjacent atom on the donor side is added to the consideration. The angle that determines when the hydrogen bond is considered to be broken is set to a default deviation of 34° [39]. Hydrogen bond acceptors are to be regarded symmetrically. 

Although some would contend that hydrogen bonding is merely an extreme manifestation of dipole-dipole interactions, it appears to be sufficiently different to warrant a short, separate discussion. In addition, there is no universal agreement on the best description of the nature of the forces in the hydrogen bond. 

Macroscopically the effects of hydrogen bonding are seen indirectly in the greatly increased melting and boiling points of such species as NH3, H,0. and HF. This 

Von der Wools distances and observed distances (pm) for some common hydrogen bonds0 

Systematic analyses of crystallographic data for hydrogen bonds have revealed a range of geometries and have led to proposals for rules to rationalize or predict hydrogen bonding patterns.111 An energetic preference for linear or near-linear 

An interesting hydrate is that of the hydronium ion in the gas phase. It consists of a dodecahedral cage of water molecules enclosing the hydronium ion HjO HjO) Each water molecule is bonded to three others in the dodecahedron (Fig. 8.8a). Of the various possible hydrates of HjO+ in the gas phase, H]O (H2O)20 is by far the most stable.21 

Von der Waals distance and obsomd dtstoncos (pm) lor som common hydrognn 

Hydrates and The hydration of ions upon solution in water has been mentioned previously and its 

If is a common error to assume that hydrogen bonding occurs in any molecule that contains hydrogen. It only occurs to a significant degree n moleoleswith N—H. O—H, orF—H bonds. 

Amino adds are the building blocks of proteins. Each amino add has both amino (—NH2) and carboxy (—COOH) functional groups, in addition to a side group that may be polar or nonpolar. Glycine, alanine, valine, and glutamic add are 4 of the 20 amino acids that make up human pro-tdns. The characteristic side group of each amino acid is shaded to identify it  

Proteins form when amino acids are joined together with peptide bonds, also known as amide linkages, in which the caiboxy group on one amino acid connects to the amine group of the next. The carboxy group of the second amino acid connects to the amine group of another, and so on. A water molecule is eliminated with the formation of each peptide bond. The side group of each amino acid, which sticks out from the backbone of the protein, is referred to as an amino acid residue, or simply as a residue  

9 Jeffrey, G. A., An Introduction to Hydrogen Bonding, Oxford University Press Oxford, 1997. 

Hydrogen bonding has tremendous effects on molecular properties. It is the strong hydrogen bonding in water that makes its boiling point of 100 °C some 160 °C higher than the heavier H2S, simply 

A—H B interaction Mainly covalent Mainly electrostatic Electrostatic 

Examples Gas phase dimers with strong acids/bases Acids Minor components of bifnrcated bonds 

Berryman, O. B., Bryantsev, V. S., Stay, D. P., Johnson, D. W. and Hay, B. P., Strnctnral criteria for the design of anion receptors The interaction of halides with electron-deficient arenes , J. Am. Chem. Soc., 2007, 129, 48-58. 

In many energetic materials, significant hydrogen bonding is observed. This is particularly true of a number of salts of the dinitramide anion. In 

Following Espinosa et al. [38], the dissociation energies for the hydrogen bonds have been estimated as  

Here v(rc) is a potential energy density at the critical point, which was calculated as in the previous section. The correlation between the length of the hydrogen bond, p(rc), V2p(rc), and dissociation energy are striking. 

Although X-ray crystallography has mainly been used for structure determination in the past, the diffraction data, especially if measured carefully and to high orders, contain information about the total electron density distribution in the crystal. This may be analyzed to provide essential information about the chemical properties of molecules, in particular the characterization of covalent and hydrogen bonds, both from the point of view of the valence electron density, the Laplacian of the density and derived energy density distribution. In addition, calculation of the molecular electrostatic potential indicates direction of chemical attack as well as how molecules can interact with their environment. 

As remarked in Section 5.1.2, the discovery of hydrogen bonding provoked ongoing controversy between proponents of a partial covalency and advocates of an electrostatic picture of H-bond formation. The former group emphasized the importance of resonance-type chemical forces of quantum-mechanical origin that could be represented as 

It is tempting to argue that in view of the close agreement between this electrostatic energy and the hydrogen bond energy, a true account has been obtained of the most important factors involved. But this is not so. 

Coulson concluded that the most important contribution to H-bonding is ionic resonance (5.29a). However, generations of empirical modelers have found it convenient to employ simple pairwise-additive Coulombic formulas with empirically fitted point charges to model H-bonds, and such empirical models have tended to encourage uncritical belief in the adequacy of a classical electrostatic picture of H-bonding. 

Problem Given the monomer natural charges (Qt) and the equilibrium Cartesian coordinates (xi, yt, z,) of the water dimer, H20- H20 

Solution Af+ouiomb = -8.2 kcalmol (The actual H-bond energy is ATTn-bond = -5.8 kcal mol-1.) 

2 Supramolecular Polymer Networks and Organogels 2.1 Hydrogen Bonding 

Other research groups have started to incorporate UPy motifs as side chains into higher molecular weight polymers via random copolymerization of UPy-functionalized alkene [93] or methacrylate monomers [94—96]. Coates and 

Hydrogen-bonding polymer networks can be even prepared in ionic liquids, as reported by Noro et al. [120]. In their work, an ABA triblock copolymer was synthesized containing end blocks that form hydrogen bonds with a specifically designed homopolymer via interactions of pyridine and hydroxy styrene. The prime role of the ionic liquid is to assure good solvent conditions over a wide range of 

This term is an explicit recognition of the importance of hydrogen bonding to molecular interactions. 

Many molecular mechanics potentials were developed at a time when it was computationally impractical to add large numbers of discrete water molecules to the calculation to simulate the effect of aqueous media. As such, techniques came into place that were intended to take into account the effect of solvent in some fashion. These techniques were difficult to justify physically but they were used nevertheless. 

The first modification is to simply scale the dielectric permittivity of free space (8 ) by a scale factor D to mediate or dampen the long range electrostatic interactions. Its value was often set to be between 1.0 and 78.0, the macroscopic value for water. A value of D=2.5, so that 8 =2.58q, was often used in early CHARMM calculations. 

For the periodic boundary conditions described below, the cutoff distance is fixed by the nearest image approximation to be less th an h alf th e sm allest box len gth. With a cutoff an y larger, m ore than nearest images would be included. 

When the cutoff is sharp, discontinuities in the forces and resultant loss of conservation of energy in molecular dynamics calculations can result. To minimize edge effects of a cutoff, often the cutoff is implemented with a switching or shifting function to allow the interactions to go smoothly to zero. 

In a similar approach, and as part of a larger study [36], hydrogen bonding between p-cresol and piperazine has been shown to occur in CDCI3 with the changes in diffusion coefficients relative to the pure, dilute materials in the same solvent (Table 95) being consistent with the 2 1 binding stoichiometry observed in the crystal structure (Fig. 9.27). 

Interaction through hydrogen bonds may be classified between dipole-dipole and covalent interactions (cf. Section 4.2). In proteins by far, most of the hydrogen bonding is between an amide and a carbonyl group -N - 0 = C -. 

however, for other reasons (such as hydrophobic interactions between R-groups) peptide units are forced into the nonaqueous interior of the protein, formation of hydrogen bonds between the peptide units is strongly promoted, thereby stabilizing secondary (and, possibly, tertiary) structures of the molecule. 

Three types of adsorption occur between the macromolecule of the soil and pesticide compounds chemical, physical, and hydrogen bonding. 

Chemical bonding is due to coulombic forces resulting in the formation of covalent bond, i.e., the sharing of a pair of electrons between the pesticide and various surface atoms of soil macromolecules. Chemical bonds probably occur less frequently than the others, as we will mention later, but once formed, they are the strongest. This explains why chemists often find it difficult to recover all of a pesticide they have added to a soil, even within a few minutes. 

The remaining intermolecular interaction is among the most poorly labeled terms in all of science. Although they are frequently called hydrogen bonds, hydrogen bonding interactions are not really bonds at all. They do not involve a chemical 

Among the many roles for hydrogen bonding perhaps the most important is the linking of two strands ofDNA into a double helix. 

Which type(s) of intermolecular forces need to be overcome to convert each of the following from liquids to gases (a) CH4, (b) CH3F, (c) CH3OH 

Because the structure is totally symmetrical, no dipole-dipole forces are possible. Because the H atoms in CH4 are not bound to N, O, or F, no hydrogen bonding is possible. Only dispersion forces need to be overcome to vaporize liquid methane. 

With only one fluorine atom present, the molecule is no longer symmetrical. The C—F bond is highly polar and will give the molecule a dipole moment. So to vaporize CH3F, dipole-dipole forces must be overcome. 

The resonance terminology and the double-headed arrow may give the impression that the structures are rapidly interconverting. This is not true. We must appreciate right from the start that resonance structures are entirely hypothetical. They are our (sometimes clumsy) attempt to write down on paper what the bonding in the molecule might be like, and they may depict only the extreme possibilities. The molecule is presumably happily going about its business in a form that we cannot easily depict. Nevertheless, resonance structures are extremely useful and do help us to explain chemical behaviour. 

Let us look again at the simple examples shown above and the consequences of our hypothetical resonance structures. 

We shall see that most of the reactions of simple carbonyl compounds, like formaldehyde, are a consequence of the presence of an electron-deficient carbon atom. This is accounted for in resonance theory by a contribution from the resonance structure with charge separation (see Section 7.1). The second example shows the so-called conjugate acid of acetone, formed to some extent by treating acetone with acid (see Section 7.1). Protonation in this way typically activates acetone towards reaction, and we 

Hydrogen bonds (H-bonds) describe the weak attraction of a hydrogen atom bonded to an electronegative atom, such as oxygen or nitrogen, to the lone pair electrons of another electronegative atom. These bonds are different in nature from the covalent bonds we have described they are considerably weaker than covalent bonds, bnt tnm ont to be snrprisingly important in chemistry and biochemistry. 

Let ns consider a molecnle possessing an 0-H a bond. This bond is polar becanse hydrogen is less electronegative than oxygen (see Section 2.7), and 

Recently much attention has been devoted to the detailed mechanism by which the class of enzymes called serine proteases work. These enzymes catalyze the ubiquitous and paramount cleavage of peptide bonds and all have the so-called catalytic triad (His-Asp-Ser) in common. A number of studies have suggested that a low barrier hydrogen bond (LBHB) is involved in the reaction mechanism as a partial proton transfer between His and Asp (N-H---0). In 

The intermolecular forces have been described as separate entities however, in the types of samples with which forensic scientists have to deal, it is rare to find only one force at work. When the intra- or intermolecular forces are discussed, we consider the interaction between molecules of identical natures (i.e., the hydrogen bonding between water molecules or dipole-dipole interactions between CO2 molecules). However, because of the complex matrices that are forensic science samples (containing many substances), we must consider all types of forces that may be present. 

If we return to the bonding of atoms and consider this in relation to the electronegativity of the atoms involved, we find that the larger the difference in the electronegativities, the stronger the bond. Polarity, on the other hand, refers to the intermolecular forces between the 6+ end of a molecule and the 6 end of another molecule of the same or a different molecule. 

What does this mean for a forensic scientist Molecules may contain polar covalent bonds and may or may not be polar the overall polarity of the molecule is determined by measuring the dipole moment. This dipole moment will depend upon the distance between the two ends (oppositely charged) of the bond and also on the overall degree of separation of charge between the atoms in the bond. If we find that a molecule contains equal polar bonds that balance each other around the central atom, we find that the molecule will be nonpolar, even though the bonds within it are polar. 

Because we now know that the polarity of a compound will affect the separation that might occur in a mixture, it follows that the solubility of those compounds in a solvent will also be affected by polarity. Not only do we 

When solvents are used in liquid chromatography, it is useful to know their polarity in relation to each other so that an appropriate choice of mobile phase or mobile phase combination might be chosen. Table 2.3 shows a list of commonly used solvents in mobile phases in order of increasing polarity. 

The CH3C=0 group (the acetyl group) is electron withdrawing. When it is substituted on the ring position opposite the carboxylic acid group, its inductive effect increases the strength of this acid as compared to benzoic acid. 

As the following examples show, the hybridization of the atom bonded to the hydrogen has a large effect on the acidity of that hydrogen  

In this series, as the hybridization changes from sp2 in ethane to sp1 in ethene and to sp in ethyne, the acidity increases and the pATa decreases. This is because of the relative stability of the unshared electrons in the conjugate bases of each of these compounds. 

Stabilization by resonance is used to explain many observations in organic chemistry. Resonance stabilization of a product can shift an equilibrium dramatically to the right, that is, to the product side of the reaction. Resonance can also lower the activation energy for a reaction, resulting in a considerable increase in reaction rate. What we learn 

Because of its sp hybridization, ethyne is a strong enough acid that its conjugate base can be readily generated by using one of the strong bases that are available in the laboratory. Calcium carbide, CaC2 a relatively stable solid, can be viewed as a dianion of ethyne  

Molecular size can be further limiting factor in oral absorption [61]. The Lipinski rule-of-5 proposes an upper MW limit of 500 as being acceptable for orally absorbed compounds [9]. Size and shape parameters are generally not measured, but rather calculated. One measured property is the so-called cross-sectional area, which is obtained from surface activity measurements [62], 

Molecular weight is often taken as the size descriptor of choice, mainly because it is easy to calculate and is generally in the chemist s mind. However, other size and shape properties are equally very simple to calculate and may offer a better guide to estimate potential for permeability. As yet, no systematic studies have been reported which investigate this in detail. Cross-sectional area (Ad, obtained from surface activity measurements) has been reported as being a useful size descriptor to discriminate compounds which can access the brain (Ad 80 A2) from those that are too large to cross the BBB [62]. Similar studies have been performed to define a cut-off for oral absorption [63]. 

Molecular size and hydrogen bonding have been unraveled as the two major components of log P or log D [52, 64, 65], In recent years it has been found that the 

Molecule a, liter2 atm/mol2 Molecule a, liter2 atm/mol2 

Both types of hydrogen bonds occur in pure liquids as well as in solutions. Many substances are associated at least partially in the vapor phase as a result of hydrogen bonding. For example, hydrogen cyanide is associated to give structures such as 

The association of acetic acid in the vapor phase occurs so that the molecular weight of the gas indicates that it exists as dimers  

Studies have indicated that the association of HF in the gas phase leads predominantly to dimers or hexamers with small amounts of tetramers. Hydrogen bonding in liquids such as sulfuric and phosphoric acids is responsible for them being viscous liquids that have high boiling points. 

Association of alcohols in the liquid state occurs with the formation of several types of species including chains, 

As noted above, more recent experimental studies of the ammonia dimer favor the modified cyclic structure 30. In light of these latest findings, new parameters have been developed [specifically, the VdW radius of H(N) was increased to 1.6 A, similar to that of H(O) and H(C)] resulting in interaction energies of —2.17 and —0.56 kcalmol-1 and N... N distances of 3.07 and 2.50 A for the linear and modified cyclic structures, respectively. 

TABLE 14. Energetic and structural parameters for three ammonia dimmers (29, 28 and 31) as calculated ab initio (6-31G + BSSE correction), by the original MM3 force field (MM3) and by MM3 augmented with a directional hydrogen bonding potential function (MM3-94)°. Reproduced by permission 

Two additional systems in which hydrogen bonds are expected to play a dominant role, ammonia-water complex and 2-aminoethanol, were calculated ab initio and by the new MM3 force field. Two ammonia-water complexes were considered, one with an N... H—O bridge (32) and the other with an O... H—N bridge (33). As expected from the relative H-donor/H-acceptor properties of nitrogen and oxygen, 32 was calculated 

These are short-range attractive interactions with a specific orientational character. There are three main molecular types (Cantor and Schimmel, 1980 Lehninger, 1982 Dickinson and McClements, 1995 Finkelstein and Ptitsyn, 2002 McClements, 2005)  

There is also an interaction between the OH group of an alcohol and the -electrons in molecules such as benzene when both are dissolved in an inert solvent. 

Hydrogen bond energies. Hydrogen bonds are frequently classified in terms of the enthalpy of formation of the bonds. Table 3.8 shows representative values for three arbitrary categories that are used to classify hydrogen bonds. 

will be solvent-dependent unless IA//3I = IA Hi + AH2, where A Hi and A H2 are the heats of solvation of the electron donor and the hydrogen compound, respectively, and AH3 is the heat of solvation of the complex. The enthalpy of the formation of the adduct in the gas phase, which would give the actual strength of the hydrogen bond, will be different from that in solution unless the solvation energies cancel. In many cases, this is questionable because of the electron donor properties of the solvent. There is no doubt that many of the hydrogen bond enthalpies reported in the literature may be in error because of this. Moreover, it has been shown that the extent of self-association of 

Effects of hydrogen bonding on physical and chemical properties. Hydrogen bonding produces many physical and chemical effects. The added intermolecular interaction often produces a drastic effect on melting and boiling points. For example, H20 boils at 100°C and H2S boils at -61 °C. BF3 is a gas (m.p. -127°C, b.p. -101 °C), whereas boric acid, B(OH)3, is a solid that decomposes at 185 °C. 

Although dimethyl ether and ethyl alcohol both have the empirical formula C2H60, ethyl alcohol boils at 78.5 °C, whereas dimethyl ether boils at -25 °C. The major portion of this difference is due to the extensive hydrogen bonding in ethyl alcohol that results in the formation of molecular aggregates that are more difficult to separate than are the polar dimethyl ether molecules that form no hydrogen bonds. 

FIGURE 3-18 Cancellation of Bond Dipoles due to Molecular Symmetry. 

FIGURE 3-20 Dimer Structures in the Gas Phase, (a) Known hydrogen-bonded structures. Rjj = hydrogen bond distance, (b) Proposed structures of the NH3 dimer and trimer. 

Nelson, Jr., G. T. Fraser, and W. Klemperer, Science, 1987, 238, 1670 M. Behrens, U. Buck, R. FrSchtenicht, andM. Hartmann, J. Chem. Phys., 1997,107, 7179 F. Huisken andT. Pertsch, Chem. Phys., 1988,126, 213. 

Another example is a theory of anesthesia by non-hydrogen bonding molecules such as cyclopropane, chloroform, and nitrous oxide, proposed by Pauling. These molecules are of a size and shape that can fit neatly into a hydrogen-bonded water structure with even larger open spaces than ordinary ice. Such structures, with molecules trapped in holes in a solid, are called clathrates. Pauling proposed that similar hydrogen-bonded microcrystals form even more readily in nerve tissue because of the presence of other solutes in the tissue. These microcrystals could then interfere with the transmission of nerve impulses. Similar structures of methane and water are believed to 

FIGURE 3-22 Hydrogen -Bonded Protein Structures, (a) A protein a helix. Peptide carbonyls and N—H hydrogens on adjacent turns of the helix are hydrogen-bonded. (From T. L. Brown and H. E. LeMay, Jr., Chemistry, the Central Science, Prentice Hall, Englewood Clifts, 

You might find it helpful to make a table of the similarities and differences between covalent and ionic compounds. 

Much more could be said about the importance of hydrogen bonding, as it is responsible for the low density of ice and for keeping proteins in shape in the cells of our bodies. A reference to hydrogen bonding occurs in the study of alcohols in Unit 7.1. 

Weak complexes of ANI with water molecules were investigated using various IR experiments264- 267. Infrared depletion spectroscopy induced by the IR multiphoton decomposition (IRMPD) processes allowed the difference between the structure of the neutral and ionized complexes to be differentiated. 

The complex formation between ANI and ammonia clusters has been investigated by using mass resolved excitation spectroscopy (MRES), hole burning spectroscopy (HB)268 and IR spectroscopy coupled to different ionization spectroscopies269-271. Rotational spectra of these complexes are not reported yet. Some ab initio calculations on both neutral and ionized complexes are available269 272-274. In this case, the most stable form of the neutral ANI-NH3 complex is a consequence of a hydrogen bond between a NH bond of aniline 

In both water and ammonia complexes, the clusters formed by interaction with the ANI molecule become invariably stronger following ionization. Thus, the hydrogen bond pattern involving amine could be modified by controlling its charge. 

It is worth noting that in the 1 1 complex of hydroxy aniline (or aminophenol) with a water molecule, the most stable cluster formed arises from an O—H—O interaction, in which water plays the role of hydrogen bond acceptor275. It is also remarkable that association of ANI with the CF3H molecule yields a weak C—H-N complex, but it induces a blue-shift of the CH stretch frequency276. The latter is blue-shifted by 30 cm-1, and the 

Conhnuous lines correspond to measured data points, dotted lines to calculated extrapolations 

0 distances. These two phenomena favour the development of stronger 0-H...0 hydrogen bonds (Fig. 7). 

According to Melzer (2000), OH absorption spectra from phlogopites with incompletely filled interlayers show a significant band at 3675 cm Since the intensity of this band correlates with the measured vacancy concentrations, this band may be assigned to the presence of interlayer vacancies and is not attributed to A1 in octahedral coordination. 

In synthetic annites two prominent bands at 3667 and 3535 cm with two shoulders 

This equation was found to give good estimates for a number of complexes, including x-complexes, but failed for complexes where the donor is a CH group. 

Khb is the equilibrium constant in carbon tetrachloride for the 1 1 complexation, and and p are the parameters for the donor and acceptor, respectively. Available data indicates that eq. 7 with slightly different coefficients also can describe complexation in 1,1,1-trichloroethane and the gas phase [67]. The 

Kenny followed the approach of Murray and Politzer and investigated correlations between the electrostatic potential and hydrogen bond basicity for a set of 23 heterocycles with nitrogen as the acceptor [69]. A very good linear relationship (R = 0.981) was found between the HF/6-31G computed Vmin and log Khb for formation of 1 1 complexes with 4-nitrophenol in 1,1,1-trichloroethane. It was pointed out that the Vmin versus log Khb relationship is significantly better than the correlations between aqueous basicity and hydrogen bond basicity (pKa versus log K) for the same types of system. The predictive capability of the relationship was further demonstrated for a set of five heterocycles in which all molecules contains two or more non-equivalent nitrogen donors. Four of the predictions are within 0.30 units of the experimental values, while the fifth, tetrazole, is overestimated by 0.51 units. 

In Table 4 are listed experimentally derived a2 values, together with HF/6-31G computed Vs,max values, for a group of 18 hydrogen bond donors of different types, including CH, NH and OH donors. There is an excellent linear 

HF/6-31G computed Vs ax 3iid experimental hydrogen bond acidities 

Following the introdnction of a directional hydrogen bonding potential function into MM3, the parameterization of the force field for the ammonia dimer was undertaken anew . Three conformers were considered, namely 28, 29 and a bifurcated structure 31, and were calcnlated ab initio at the 6-31G level. The results (after corrections for Basis Set Superimposition Error BSSE) favor the linear dimer over the cyclic one by 0.4 kcalmol and yield dimerization energies of —2.49, —2.09 and —0.62 kcalmol for 28, 29 and 31, respectively. A comparison of force field (original MM3 and MM3 with the directional hydrogen bonding function) and ab initio resnlts for the three ammonia 

FIGURE 3.27 The open structure of ice. (Brown Lemay, Chemistry Central Science, 4th Ed., 1988, pp. 628,946. Reprinted and Electronically reproduced by permission of Pearson Education Inc, Upper Saddle River, NJ 07458.) The rectangular lines are included to aid visualization all bonding is between hydrogen and oxygen atoms. 

More specific interactions involving the sharing of electron pairs between molecules are discussed in connection with acid-base chemistry in Chapter 6. 

Gillespie, J. Chem. Educ., 1970, 47, 18 and R. J. Gillespie, Coord. Chem. Rev., 2008, 252, 1315. The last reference provides a useful synopsis of 50 years of the VSEPR model. 

Barrow, Physical Chemistry, 6th ed., McGraw-Hill, New York, 1988, pp. 567-699 R. S. Drago, Physical Methods for Chemists, 2nd ed., Saunders College Publishing, Philadelphia, 1977, pp. 689-711. 

Why is the boiling point of SiH4 higher than that of CH4  

Covalent bond, iMframolecular Hydrogen bond, iMfermolecular 

Hydrogen bonding can occur when an H atom is bonded to an N, 0, or F atom. 

Solvent Dn.i approximate DNsbCis (calc.) DNsbOi (found) 

Consider the physical properties of the two isomeric compounds, ethanol and dimethyl etiher, presented in Table 7.3. Since the dipole moment values are so close, how can we explain the much higher normal boiling point temperature of ethanol Or, the more than double amount of energy needed to vaporize one mole of ethanol as compared to one of dimethyl ether In other words, what kind of intermolecular forces are responsible for these differences  

To answer these questions we notice that, as discussed in Section 7.4.1, oxygen in the hydroxyl group has the tendency to steal the electron from the hydrogen atom rather than share it, along with one of its own, with the latter. 

As a result, we have an oxygen atom that has more than its share of electrons and, consequently, can act as an electron donor or a proton acceptor, and a hydrogen atom, that has less than its share of electrons and can act as an electron acceptor or a proton donor. 

Molecules, thus, containing a hydrogen atom linked to an electronegative one - such as acids, alcohols, and amines - have a strong tendency to form aggregates with each other or with other molecules containing accessible electronegative atoms. 

When the aggregates formed contain molecules of the same type, such as in the aforementioned case of ethanol, the phenomenon is referred to as association, when they contain molecules of different type, such as ethanol and water ones for example, as solvation. 

When methane or fluorine gases are chilled, it is van der Waals forces that hold the molecules together in the resulting solids and liquids. They are weak, and a measure of their weakness is the low boiling temperature of liquid methane or liquid fluorine (-162 °C and -188 °C), respectively. 

Which three hydrides seem to have anomalous values  

The boiling temperatures of the noble gases (Group VIII) and of the hydrides of Groups IV-VII. 

How many other water molecules is each water molecule surrounded by  

Look at one of the two water molecules which is located wholly within the unit cell each is tetrahedrally surrounded by only four others. 

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Bfi and 022- However, in the second binary, intermolecular forces between unlike molecules are much stronger than those between like molecules chloroform and ethyl acetate can strongly hydrogen bond with each other but only very weakly with them-... 

The virial equation is appropriate for describing deviations from ideality in those systems where moderate attractive forces yield fugacity coefficients not far removed from unity. The systems shown in Figures 2, 3, and 4 are of this type. However, in systems containing carboxylic acids, there prevails an entirely different physical situation since two acid molecules tend to form a pair of stable hydrogen bonds, large negative... 

Moderate errors in the total pressure calculations occur for the systems chloroform-ethanol-n-heptane and chloroform-acetone-methanol. Here strong hydrogen bonding between chloroform and alcohol creates unusual deviations from ideality for both alcohol-chloroform systems, the activity coefficients show... 

Breaking one or more carbon-hydrogen bonds results in the following ... 

The different kinds of intermolecular forces (dispersion, dipole-dipole, hydrogen bonding, etc. see Section VI-1) may not equally contribute to A-A, B-B, and A-B... 

Figure IV-10 illustrates how F may vary with film pressure in a very complicated way although the v-a plots are relatively unstructured. The results correlated more with variations in film elasticity than with its viscosity and were explained qualitatively in terms of successive film structures with varying degrees of hydrogen bonding to the water substrate and varying degrees of structural regularity. Note the sensitivity of k to frequency a detailed study of the dispersion of k should give information about the characteristic relaxation times of various film structures.

In order to include other interactions such as dipolar or hydrogen bonding, many semiempirical approaches have been tried [196, 197, 200], including adding terms to Eq. X-45 [198, 201] or modifying the definition of [202, 199]. Perhaps the most well-known of these approaches comes from Fowkes [203, 204] suggestion that the interactions across a water-hydrocarbon interface are dominated by dispersion forces such that Eq. X-45 could be modified as... 

There is a fair amount of work reported with films at the mercury-air interface. Rice and co-workers [107] used grazing incidence x-ray diffraction to determine that a crystalline stearic acid monolayer induces order in the Hg substrate. Quinone derivatives spread at the mercury-n-hexane interface form crystalline structures governed primarily by hydrogen bonding interactions [108]. 

Gavezzotti A and Filippini G 1994 Geometry of the intermolecular XH.. . Y (X,Y = N,0) hydrogen bond and the calibration of empirical hydrogen-bond potentials J. Phys. Chem. 98 4831... 

Scheiner S 1997 Hydrogen Bonding A Theoretical Perspective (New York Oxford) A survey of research on hydrogen bonding with emphasis on tiieoretical calculations. 1994 van der Waals molecules Chem. Rev. 94 1721... 

Figure Al.7.14. 3.4 mn x 3.4 mn STM images of 1-docosanol physisorbed onto a graphite surface in solution. This image reveals the hydrogen-bonding alcohol molecules assembled in lamellar fashion at the liquid-solid interface. Each bright circular region is attributed to the location of an individual hydrogen...

Gragson D E and Richmond G I 1998 Investigations of the structure and hydrogen bonding of water molecules at liquid surfaces by vibrational sum frequency spectroscopy J. Phys. Chem. 102 3847... 

This arises because as the temperature in increased from ambient, the main initial effect is to loosen the hydrogen-bonded local stmcture that iitiribits reorientation. Flowever, at higher temperatures, the themial motion of the water molecules becomes so marked that cluster fomration becomes iitiiibited. 

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