Hydrogen Bond - Wikipedia

 

 

Definitions and general characteristics

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In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named the proton acceptor, whereas the one covalently bound to the hydrogen is named the proton donor. This nomenclature is recommended by the IUPAC.[6] The hydrogen of the donor is protic and therefore can act as a Lewis acid and the acceptor is the Lewis base. Hydrogen bonds are represented as H···Y system, where the dots represent the hydrogen bond. Liquids that display hydrogen bonding (such as water) are called associated liquids.[citation needed]

 

 

Hydrogen bonds arise from a combination of electrostatics (multipole-multipole and multipole-induced multipole interactions), covalency (charge transfer by orbital overlap), and dispersion (London forces).[6]

In weaker hydrogen bonds,[13] hydrogen atoms tend to bond to elements such as sulfur (S) or chlorine (Cl); even carbon (C) can serve as a donor, particularly when the carbon or one of its neighbors is electronegative (e.g., in chloroform, aldehydes and terminal acetylenes).[14][15] Gradually, it was recognized that there are many examples of weaker hydrogen bonding involving donor other than N, O, or F and/or acceptor Ac with electronegativity approaching that of hydrogen (rather than being much more electronegative). Although weak (about 4.2 kJ/mol (1 kcal/mol)), "non-traditional" hydrogen bonding interactions are ubiquitous and influence structures of many kinds of materials.[citation needed]

The definition of hydrogen bonding has gradually broadened over time to include these weaker attractive interactions. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. This definition specifies:

The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X−H in which X is more electronegative than H, and an atom or a group of atoms in the same or another molecule, in which there is evidence of bond formation.[16]

Bond strength

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Hydrogen bonds can vary in strength from weak (1–2 kJ/mol) to strong (161.5 kJ/mol in the bifluoride ion, HF2-).[17][18] Typical enthalpies in vapor include:[19]

• F−H···F− (161.5 kJ/mol (38.6 kcal/mol)), illustrated uniquely by HF2-
• O−H···N (29 kJ/mol (6.9 kcal/mol)), illustrated water-ammonia
• O−H···O (21 kJ/mol (5.0 kcal/mol)), illustrated water-water, alcohol-alcohol
• N−H···N (13 kJ/mol (3.1 kcal/mol)), illustrated by ammonia-ammonia
• N−H···O (8 kJ/mol (1.9 kcal/mol)), illustrated water-amide
• OH+3···OH2 (18 kJ/mol (4.3 kcal/mol))[20]

The strength of intermolecular hydrogen bonds is most often evaluated by measurements of equilibria between molecules containing donor and/or acceptor units, most often in solution.[21] The strength of intramolecular hydrogen bonds can be studied with equilibria between conformers with and without hydrogen bonds. The most important method for the identification of hydrogen bonds also in complicated molecules is crystallography, sometimes also NMR-spectroscopy. Structural details, in particular distances between donor and acceptor which are smaller than the sum of the van der Waals radii can be taken as indication of the hydrogen bond strength. One scheme gives the following somewhat arbitrary classification: those that are 63 to 167 kJ/mol (15 to 40 kcal/mol), 21 to 63 kJ/mol (5 to 15 kcal/mol), and 0 to 21 kJ/mol (0 to 5 kcal/mol) are considered strong, moderate, and weak, respectively.[18]

Hydrogen bonds involving C−H bonds are both very rare and weak.[22]

Resonance assisted hydrogen bond

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The resonance assisted hydrogen bond (commonly abbreviated as RAHB) is a strong type of hydrogen bond. It is characterized by the π-delocalization that involves the hydrogen and cannot be properly described by the electrostatic model alone. This description of the hydrogen bond has been proposed to describe unusually short distances generally observed between O=C−OH··· or ···O=C−C=C−OH.[23]

Structural details

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The X−H distance is typically ≈110 pm, whereas the H···Y distance is ≈160 to 200 pm. The typical length of a hydrogen bond in water is 197 pm. The ideal bond angle depends on the nature of the hydrogen bond donor. The following hydrogen bond angles between a hydrofluoric acid donor and various acceptors have been determined experimentally:[24]

Acceptor···donor	VSEPR geometry	Angle
HCN···HF	C≡N···H angle: linear	180°
H2CO···HF	C=O···H angle: trigonal planar	120°
H2O···HF	H-O···H angle: pyramidal	46°
H2S···HF	H-S···H angle: pyramidal	89°
SO2···HF	S=O···H angle: trigonal	145°

Spectroscopy

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Strong hydrogen bonds are revealed by downfield shifts in the 1H NMR spectrum. For example, the acidic proton in the enol tautomer of acetylacetone appears at ⁠ δ H {\displaystyle \delta _{\text{H}}}  ⁠ 15.5, which is about 10 ppm downfield of a conventional alcohol.[25]

In the IR spectrum, hydrogen bonding shifts the X−H stretching frequency to lower energy (i.e. the vibration frequency decreases). This shift reflects a weakening of the X−H bond. Certain hydrogen bonds - improper hydrogen bonds - show a blue shift of the X−H stretching frequency and a decrease in the bond length.[26] H-bonds can also be measured by IR vibrational mode shifts of the acceptor. The amide I mode of backbone carbonyls in α-helices shifts to lower frequencies when they form H-bonds with side-chain hydroxyl groups.[27] The dynamics of hydrogen bond structures in water can be probed by this OH stretching vibration.[28] In the hydrogen bonding network in protic organic ionic plastic crystals (POIPCs), which are a type of phase change material exhibiting solid-solid phase transitions prior to melting, variable-temperature infrared spectroscopy can reveal the temperature dependence of hydrogen bonds and the dynamics of both the anions and the cations.[29] The sudden weakening of hydrogen bonds during the solid-solid phase transition seems to be coupled with the onset of orientational or rotational disorder of the ions.[29]

Theoretical considerations

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Hydrogen bonding is of persistent theoretical interest.[30] According to a modern description O:H−O integrates both the intermolecular O:H lone pair ":" nonbond and the intramolecular H−O polar-covalent bond associated with O−O repulsive coupling.[31]

Quantum chemical calculations of the relevant interresidue potential constants (compliance constants) revealed[how?] large differences between individual H bonds of the same type. For example, the central interresidue N−H···N hydrogen bond between guanine and cytosine is much stronger in comparison to the N−H···N bond between the adenine-thymine pair.[32]

Theoretically, the bond strength of the hydrogen bonds can be assessed using NCI index, non-covalent interactions index, which allows a visualization of these non-covalent interactions, as its name indicates, using the electron density of the system.[citation needed]

Interpretations of the anisotropies in the Compton profile of ordinary ice claim that the hydrogen bond is partly covalent.[33] However, this interpretation was challenged [34] and subsequently clarified.[35]

Most generally, the hydrogen bond can be viewed as a metric-dependent electrostatic scalar field between two or more intermolecular bonds. This is slightly different from the intramolecular bound states of, for example, covalent or ionic bonds. However, hydrogen bonding is generally still a bound state phenomenon, since the interaction energy has a net negative sum. The initial theory of hydrogen bonding proposed by Linus Pauling suggested that the hydrogen bonds had a partial covalent nature. This interpretation remained controversial until NMR techniques demonstrated information transfer between hydrogen-bonded nuclei, a feat that would only be possible if the hydrogen bond contained some covalent character.[36]