How strong is a hydrogen bond?
Hydrogen bond
Hydrogen bonds can vary in strength from weak (1–2 kJ mol) to strong (161.5 kJ mol in the ion ). Typical enthalpies in vapor include:
Self-healing hydrogels
Hydrogen bonding is a strong intermolecular force that forms a special type of dipole-dipole attraction. Hydrogen bonds form when a hydrogen atom bonded to a strongly electronegative atom is around another electronegative atom with a lone pair of electrons. Hydrogen bonds are stronger than normal dipole-dipole interactions and dispersion forces but they remain weaker than covalent and ionic bonds. In hydrogels, structure and stability of water molecules are highly affected by the bonds. The polar groups in the polymer strongly bind water molecules and form hydrogen bonds which also cause hydrophobic effects to occur. These hydrophobic effects can be exploited to design chemically crosslinked hydrogels that exhibit self healing abilities. The hydrophobic effects combined with the hydrophilic effects within the hydrogel structure can be balanced through dangling side chains that mediates the hydrogen bonding that occurs between two separate hydrogel pieces or across a ruptured hydrogel.
A dangling side chain is a hydrocarbon chain side chains that branch off of the backbone of the polymer. Attached to the side chain are polar functional groups. The side chains "dangle" across the surface of the hydrogel, allowing it to interact with other functional groups and form new bonds. The ideal side chain would be long and flexible so it could reach across the surface to react, but short enough to minimize steric hindrance and collapse from the hydrophobic effect. The side chains need to keep both the hydrophobic and hydrophilic effects in balance. In a study performed by the University of California San Diego to compare healing ability, hydrogels of varying side chain lengths with similar crosslinking contents were compared and the results showed that healing ability of the hydrogels depends nonmonotonically on the side chain length. With shorter side chain lengths, there is limited reach of the carboxyl group which decreases the mediation of the hydrogen bonds across the interface. As the chain increases in length, the reach of the carboxyl group becomes more flexible and the hydrogen bonds can mediated. However, when a side chain length is too long, the interruption between the interaction of the carboxyl and amide groups that help to mediate the hydrogen bonds. It can also accumulate and collapse the hydrogel and prevent the healing from occurring.
Hydrogen bond
Hydrogen bonds are important in the structure of cellulose and derived polymers in its many different forms in nature, such as cotton and flax.
Hydrogen bond
Where the bond strengths are more equivalent, one might instead find the atoms of two interacting water molecules partitioned into two polyatomic ions of opposite charge, specifically hydroxide (OH) and hydronium (HO). (Hydronium ions are also known as "hydroxonium" ions.)
Hydrogen-bond catalysis
In general, acidity of donor sites correlates well with the strength of the donor. For example, it is a common strategy to add electron-withdrawing aryl substituents on a thiourea catalyst, which can increase its acidity and thus the strength of its hydrogen bonding. However, it is still unclear how donor strength correlates with desired reactivity. Importantly, more acidic catalysts are not necessarily more effective. For instance, ureas are less acidic than thioureas by roughly 6 pKa units, but it is not generally true that ureas are significantly worse are catalyzing reactions.
Hydrogen embrittlement
As the strength of steels increases, the susceptibility to hydrogen embrittlement increases. In high-strength steels, anything above a hardness of HRC 32 may be susceptible to early hydrogen cracking after plating processes that introduce hydrogen. They may also experience long-term failures anytime from weeks to decades after being placed in service due to accumulation of hydrogen over time from cathodic protection and other sources. Numerous failures have been reported in the hardness range from HRC 32-36 and more above; therefore, parts in this range should be checked during quality control to ensure they are not susceptible.
Hydrogen bond
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.
Hydrogen bond
Several studies have shown that hydrogen bonds play an important role for the stability between subunits in multimeric proteins. For example, a study of sorbitol dehydrogenase displayed an important hydrogen bonding network which stabilizes the tetrameric quaternary structure within the mammalian sorbitol dehydrogenase protein family.
Chlorine
At room temperature, hydrogen chloride is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the larger electronegative chlorine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen chloride at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised. Hydrochloric acid is a strong acid (p"K" = −7) because the hydrogen bonds to chlorine are too weak to inhibit dissociation. The HCl/HO system has many hydrates HCl·"n"HO for "n" = 1, 2, 3, 4, and 6. Beyond a 1:1 mixture of HCl and HO, the system separates completely into two separate liquid phases. Hydrochloric acid forms an azeotrope with boiling point 108.58 °C at 20.22 g HCl per 100 g solution; thus hydrochloric acid cannot be concentrated beyond this point by distillation.
Hydrogen fluoride
Although a diatomic molecule, HF forms relatively strong intermolecular hydrogen bonds. Solid HF consists of zigzag chains of HF molecules. The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm. Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.
Bond valence method
A chemical structure can be represented by a bond network of the kind familiar in molecular diagrams. The infinitely connected bond networks found in crystals can be simplified into finite networks by extracting one formula unit and reconnecting any broken bonds to each other. If the bond network is not known, a plausible network can be created by connecting well matched cations and anions that satisfy Eq. 4. If the finite network contains only cation-anion bonds, every bond can be treated as an electric capacitor (two equal and opposite charges linked by electrostatic flux). The bond network is thus equivalent to a capacitive electrical circuit with the charge on each capacitor being equivalent to the bond valence. The individual bond capacitors are not initially known, but in the absence of any information to the contrary we assume that they are all equal. In this case the circuit can be solved using the Kirchhoff equations, yielding the valences of each bond. Eq. 2 can then be used to calculate bond lengths which are found to lie within a few picometres of the observed bond lengths if no additional constraints are present. Additional constraints include electronic anisotropies (lone pairs and Jahn-Teller distortions) or steric constraints, (bonds stretched or compressed in order to fit them into three-dimensional space). Hydrogen bonds are an example of a steric constraint. The repulsion resulting from the close approach of the donor and acceptor atoms causes the bonds to be stretched, and under this constraint the distortion theorem predicts that the hydrogen atom will move off-center.
Hydrogen bond
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:
Strong hydrogen bonds are revealed by downfield shifts in the H NMR spectrum. For example, the acidic proton in the enol tautomer of acetylacetone appears at δ15.5, which is about 10 ppm downfield of a conventional alcohol.
Hydrogel encapsulation of quantum dots
Increasing the PEG length increases the solubility of the polymer. However, if the PEG chain is too long the micelle will become unstable. It has been observed that a stable hydrogel can only be formed with PEG backbones weighing between six and ten thousand Daltons.
Carbon monoxide
The bond length between the carbon atom and the oxygen atom is 112.8 pm. This bond length is consistent with a triple bond, as in molecular nitrogen (N), which has a similar bond length (109.76 pm) and nearly the same molecular mass. Carbon–oxygen double bonds are significantly longer, 120.8 pm in formaldehyde, for example. The boiling point (82 K) and melting point (68 K) are very similar to those of N (77 K and 63 K, respectively). The bond-dissociation energy of 1072 kJ/mol is stronger than that of N (942 kJ/mol) and represents the strongest chemical bond known.
Chalcogen bond
Due to a chalcogen's ability to function as an electron donor, many systems will balance between hydrogen bonding with the chalcogen as a donor or chalcogen bonding. This balance can be observed in a series of self-bonding intermolecular interactions between various substituted chalcogens. In cases with hard chalcogen atoms as acceptors, the balance favors H-bonding between chalcogen and methyl H. However, as the acceptor atoms move down the group, chalcogen-chalcogen bonding will be favored. It is hypothesized that electrostatic forces should only dominate in a chalcogen-chalcogen bond with lighter congeners, and instead that dispersion forces dominate in the cases of heavier congeners
Single bond
Usually, a single bond is a sigma bond. An exception is the bond in diboron, which is a pi bond. In contrast, the double bond consists of one sigma bond and one pi bond, and a triple bond consists of one sigma bond and two pi bonds (Moore, Stanitski, and Jurs 396). The number of component bonds is what determines the strength disparity. It stands to reason that the single bond is the weakest of the three because it consists of only a sigma bond, and the double bond or triple bond consist not only of this type of component bond but also at least one additional bond.
Transition metal hydride
Although these properties are interrelated, they are not interdependent. A metal hydride can be a thermodynamically a weak acid and a weak H donor; it could also be strong in one category but not the other or strong in both. The H strength of a hydride also known as its hydride donor ability or hydricity corresponds to the hydride's Lewis base strength. Not all hydrides are powerful Lewis bases. The base strength of hydrides vary as much as the pKa of protons. This hydricity can be measured by heterolytic cleaving hydrogen between a metal complex and base with a known pK then measuring the resulting equilibrium. This presupposes that the hydride doesn't heterolytically or homolytically react with itself to reform hydrogen. A complex would homolytically react with itself if the homolytic M-H bond is worth less than half of the homolytic H-H bond. Even if the homolytic bond strength above that threshold the complex is still susceptible to radical reaction pathways.
Direct bonding
From room temperature to 150 °C no important interface reactions occur and the surface energy is stable. Between 150 °C and 300 °C more Si-F-H-Si bonds are formed. Above 300 °C the desorption of hydrogen and fluoride from the wafer surface leads to redundant hydrogen atoms that diffuse in the silicon crystal lattice or along interface. As a result, covalent Si-Si bonds start to establish between opposing surfaces. At 700 °C the transition to Si-Si bonds is completed.
The bonding energy reaches cohesive strengths of bulk silicon (compare diagram of surface energy).
Resonance (chemistry)
Hydrogenation of one mole of double bonds delivers 119.7 kJ (28.6 kcal), as can be deduced from the last step, the hydrogenation of cyclohexene. In benzene, however, 23.4 kJ (5.6 kcal) are needed to hydrogenate one mole of double bonds. The difference, being 143.1 kJ (34.2 kcal), is the empirical resonance energy of benzene. Because 1,3-cyclohexadiene also has a small delocalization energy (7.6 kJ or 1.8 kcal/mol) the net resonance energy, relative to the localized cyclohexatriene, is a bit higher: 151 kJ or 36 kcal/mol.
This measured resonance energy is also the difference between the hydrogenation energy of three 'non-resonance' double bonds and the measured hydrogenation energy:
Halogen
All halogens form binary compounds with hydrogen known as the hydrogen halides: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAt). All of these compounds form acids when mixed with water. Hydrogen fluoride is the only hydrogen halide that forms hydrogen bonds. Hydrochloric acid, hydrobromic acid, hydroiodic acid, and hydroastatic acid are all strong acids, but hydrofluoric acid is a weak acid.
DNA
The two types of base pairs form different numbers of hydrogen bonds, AT forming two hydrogen bonds, and GC forming three hydrogen bonds (see figures, right).
DNA with high GC-content is more stable than DNA with low GC-content.
As noted above, most DNA molecules are actually two polymer strands, bound together in a helical fashion by noncovalent bonds; this double-stranded (dsDNA) structure is maintained largely by the intrastrand base stacking interactions, which are strongest for G,C stacks. The two strands can come apart—a process known as melting—to form two single-stranded DNA (ssDNA) molecules. Melting occurs at high temperature, low salt and high pH (low pH also melts DNA, but since DNA is unstable due to acid depurination, low pH is rarely used).
Intermolecular force
A "hydrogen bond" is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is bonded to either nitrogen, oxygen, or fluorine. The hydrogen bond is often described as a strong electrostatic dipole-dipole interaction. However, it also has some features of covalent bonding: it is directional, stronger than a van der Waals force interaction, produces interatomic distances shorter than the sum of van der Waals radius, and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence.
Hydrogen bond
Hydrogen bonding is a key to the design of drugs. According to Lipinski's rule of five the majority of orally active drug tend to have between five and ten hydrogen bonds. These interactions exist between nitrogen–hydrogen and oxygen–hydrogen centers. As with many other rules of thumb, many exceptions exist.
Solvation
Hydrogen bonding among solvent and solute molecules depends on the ability of each to accept H-bonds, donate H-bonds, or both. Solvents that can donate H-bonds are referred to as protic, while solvents that do not contain a polarized bond to a hydrogen atom and cannot donate a hydrogen bond are called aprotic. H-bond donor ability is classified on a scale (α). Protic solvents can solvate solutes that can accept hydrogen bonds. Similarly, solvents that can accept a hydrogen bond can solvate H-bond-donating solutes. The hydrogen bond acceptor ability of a solvent is classified on a scale (β). Solvents such as water can both donate and accept hydrogen bonds, making them excellent at solvating solutes that can donate or accept (or both) H-bonds.
Bent bond
In these compounds, it is not possible for the carbon atoms to assume the 109.5° bond angles with standard sp hybridization. Increasing the p-character to sp (i.e. s-density and p-density) makes it possible to reduce the bond angles to 60°. At the same time, the carbon-to-hydrogen bonds gain more s-character, which shortens them. In cyclopropane, the maximum electron density between two carbon atoms does not correspond to the internuclear axis, hence the name "bent bond". In cyclopropane, the "interorbital angle" is 104°. This bending can be observed experimentally by X-ray diffraction of certain cyclopropane derivatives: the deformation density is outside the line of centers between the two carbon atoms. The carbon–carbon bond lengths are shorter than in a regular alkane bond: 151 pm versus 153 pm.
Nucleic acid secondary structure
In molecular biology, two nucleotides on opposite complementary DNA or RNA strands that are connected via hydrogen bonds are called a base pair (often abbreviated bp). In the canonical Watson-Crick base pairing, adenine (A) forms a base pair with thymine (T) and guanine (G) forms one with cytosine (C) in DNA. In RNA, thymine is replaced by uracil (U). Alternate hydrogen bonding patterns, such as the wobble base pair and Hoogsteen base pair, also occur—particularly in RNA—giving rise to complex and functional tertiary structures. Importantly, pairing is the mechanism by which codons on messenger RNA molecules are recognized by anticodons on transfer RNA during protein translation. Some DNA- or RNA-binding enzymes can recognize specific base pairing patterns that identify particular regulatory regions of genes.
Hydrogen bonding is the chemical mechanism that underlies the base-pairing rules described above. Appropriate geometrical correspondence of hydrogen bond donors and acceptors allows only the "right" pairs to form stably. DNA with high GC-content is more stable than DNA with low GC-content, but contrary to popular belief, the hydrogen bonds do not stabilize the DNA significantly and stabilization is mainly due to stacking interactions.
Conjugate acid
We say that the water molecule is the conjugate acid of the hydroxide ion after the latter received the hydrogen proton donated by ammonium. On the other hand, ammonia is the conjugate base for the acid ammonium after ammonium has donated a hydrogen ion towards the production of the water molecule. We can also refer to OH- as a conjugate base of , since the water molecule donates a proton towards the production of in the reverse reaction, which is the predominating process in nature due to the strength of the base over the hydroxide ion. Based on this information, it is clear that the terms "Acid", "Base", "conjugate acid", and "conjugate base" are not fixed for a certain chemical species; but are interchangeable according to the reaction taking place.
The strength of a conjugate acid is directly proportional to its dissociation constant. If a conjugate acid is strong, its dissociation will have a higher equilibrium constant and the products of the reaction will be favored. The strength of a conjugate base can be seen as the tendency of the species to "pull" hydrogen protons towards itself. If a conjugate base is classified as strong, it will "hold on" to the hydrogen proton when in solution and its acid will not dissociate. 
On the other hand, if a species is classified as a strong acid, its conjugate base will be weak in nature. An example of this case would be the dissociation of Hydrochloric acid in water. Since HCl is a strong acid (it dissociates to a great extent), its conjugate base () will be a weak conjugate base. Therefore, in this system, most will be in the form of a Hydronium ion instead of attached to a Cl anion and the conjugate base will be weaker than a water molecule.
Hydrogen isotope biogeochemistry
Hydrogen atoms easily separate from electronegative bonds such as hydroxyl bonds (O-H), nitrogen bonds (N-H), and thiol/mercapto bonds (S-H) on hour to day long timescales. This rapid exchange is particularly problematic for measurements of bulk organic material with these functional groups because isotope compositions are more likely to reflect the source water and not the isotope effect. For this reason, records of paleoclimate that are not measuring ancient waters, rely on other isotopic markers. Advancements in the 1990s held promising potential to resolve this problem: samples were equilibrated with two variations of heavy water and compared. Their ratios represent an exchange factor that can calibrate measurements to correct for hydrogen and deuterium swapping.
Hydrogen bond
A hydrogen bond is a partially electrostatic attraction between a hydrogen (H) atom which is bound to a more electronegative atom or group, such as nitrogen (N), oxygen (O), or fluorine (F)—the hydrogen bond donor—and another adjacent atom bearing a lone pair of electrons—the hydrogen bond acceptor.
Hydrogen-bridged cations
The most significant trait of bridging hydrogens in NMR spectra is their upfield shift. Due to the increased number of bonds that the hydrogen atom is part of, the higher electron density around hydrogen shields the H nucleus, causing its chemical shift to appear at negative ppm. Characterization of the dimethylcyclodecyl ion by Sorensen in 1978 relied on low-temperature NMR, as the ion slowly decomposes at above -70 °C, and showed the bridging hydrogen at δ -3.9. McMurry's characterization of the more stable "in"-bicyclo[4,4,4]-tetradecyl cation showed two broad 12 proton peaks at δ 2.5 and 1.9, and a broad proton singlet at δ -3.46. Depending on the length of bridging chains in a similar bridged bicyclo[3,3,1]-nonyl cations the upfield shift of the bridging hydrogen can vary from δ -6.50 to 0.07. This is actually due to a change in C-H-C bond angle, where change in bond angle affects the bonding nature of the 3 center-2 electron bond.
Carbon monoxide
The carbon monoxide has a very high bond-dissociation energy, the strongest of any neutral molecule, 11.65 eV. Carbon and oxygen together have a total of 10 electrons in the valence shell. Following the octet rule for both carbon and oxygen, the two atoms form a triple bond, with six shared electrons in three bonding molecular orbitals, rather than the usual double bond found in organic carbonyl compounds. Since four of the shared electrons come from the oxygen atom and only two from carbon, one bonding orbital is occupied by two electrons from oxygen, forming a dative or dipolar bond. This causes a C←O polarization of the molecule, with a small negative charge on carbon and a small positive charge on oxygen. The other two bonding orbitals are each occupied by one electron from carbon and one from oxygen, forming (polar) covalent bonds with a reverse C→O polarization, since oxygen is more electronegative than carbon. In the free carbon monoxide, a net negative charge δ remains at the carbon end and the molecule has a small dipole moment of 0.122 D.
Supramolecular assembly
Hydrogen bond play an essential role in the assembly of secondary and tertiary structures of large biomolecules. DNA double helix is formed by hydrogen bonding between nucleobases: adenine and thymine forms two hydrogen bonds, while guanine and cytosine forms three hydrogen bonds (Figure "Hydrogen bonds in (a) DNA duplex formation"). Another prominent example of hydrogen bond-assisted assembly in nature is the formation of protein secondary structures. Both the α-helix and β-sheet are formed through hydrogen bonding between the amide hydrogen and the amide carbonyl oxygen (Figure "Hydrogen bonds in (b) protein β-sheet structure").
Pi bond
A pi bond is weaker than a sigma bond, but the combination of pi and sigma bond is stronger than either bond by itself. The enhanced strength of a multiple bond versus a single (sigma bond) is indicated in many ways, but most obviously by a contraction in bond lengths. For example, in organic chemistry, carbon–carbon bond lengths are about 154 pm in ethane, 134 pm in ethylene and 120 pm in acetylene. More bonds make the total bond shorter and stronger.
Hydrogen-bond catalysis
A notable example of the anion-binding mechanism is the hydrocyanation of imines catalyzed by Jacobsen's amido-thiourea catalyst depicted in the below diagram. This reaction is also one of the most extensively studied through computational, spectroscopic, labeling and kinetic experiments. While direct addition of cyanide to a catalyst-bound imine was considered, an alternative mechanism involving formation of an iminium-cyanide ion pair controlled by catalyst was calculated to have a barrier that is lower by 20 kcal/mol. The proposed most likely mechanism begins with binding of the catalyst to HNC, which exists in equilibrium with HCN. This complex then protonates a molecule of imine, forming an iminium-cyanide ion pair with the catalyst binding and stabilizing the cyanide anion. The iminium is thought to also interact with the amide carbonyl on the catalyst molecule (see bifunctional catalysis below). The bound cyanide anion then rotates, and attacks the iminium through carbon. The investigators conclude that though imine-urea binding was observed through spectroscopy and was supported by early kinetic experiments, imine binding is off-cycle and all evidence points toward this mechanism involving thiourea-bound cyanide.
It is often difficult to distinguish between hydrogen-bond catalysis and general acid catalysis. Hydrogen-bond donors can have varying acidity, from mild to essentially strong Brønsted acids like phosphoric acids. Looking at the extent of proton transfer over the course of the reaction is challenging and has not been investigated thoroughly in most reactions. Nevertheless, strong acid catalysts are often grouped with hydrogen-bond catalysts as they represent an extreme on this continuum and their catalytic behaviors share similarities. The mechanism of activation for these reactions involves initial protonation of the electrophilic partner. This has the effect of rendering the substrate more electrophilic and creating an ion pair, through which it is possible to transfer stereochemical information.
Hydrogen bond
In these macromolecules, bonding between parts of the same macromolecule cause it to fold into a specific shape, which helps determine the molecule's physiological or biochemical role. For example, the double helical structure of DNA is due largely to hydrogen bonding between its base pairs (as well as pi stacking interactions), which link one complementary strand to the other and enable replication.
Hydrogen bond
The hydrogen-bond networks make both natural and synthetic polymers sensitive to humidity levels in the atmosphere because water molecules can diffuse into the surface and disrupt the network. Some polymers are more sensitive than others. Thus nylons are more sensitive than aramids, and nylon 6 more sensitive than nylon-11.
Alkali metal
The 1s electron configuration of hydrogen, while superficially similar to that of the alkali metals (ns), is unique because there is no 1p subshell. Hence it can lose an electron to form the hydron H, or gain one to form the hydride ion H. In the former case it resembles superficially the alkali metals; in the latter case, the halogens, but the differences due to the lack of a 1p subshell are important enough that neither group fits the properties of hydrogen well. Group 14 is also a good fit in terms of thermodynamic properties such as ionisation energy and electron affinity, but makes chemical nonsense because hydrogen cannot be tetravalent. Thus none of the three placements are entirely satisfactory, although group 1 is the most common placement (if one is chosen) because the hydron is by far the most important of all monatomic hydrogen species, being the foundation of acid-base chemistry. As an example of hydrogen's unorthodox properties stemming from its unusual electron configuration and small size, the hydrogen ion is very small (radius around 150 fm compared to the 50–220 pm size of most other atoms and ions) and so is nonexistent in condensed systems other than in association with other atoms or molecules. Indeed, transferring of protons between chemicals is the basis of acid-base chemistry. Also unique is hydrogen's ability to form hydrogen bonds, which are an effect of charge-transfer, electrostatic, and electron correlative contributing phenomena. While analogous lithium bonds are also known, they are mostly electrostatic. Nevertheless, hydrogen can take on the same structural role as the alkali metals in some molecular crystals, and has a close relationship with the lightest alkali metals (especially lithium).
Diborane
Diborane adopts a D structure containing four terminal and two bridging hydrogen atoms. The model determined by molecular orbital theory indicates that the bonds between boron and the terminal hydrogen atoms are conventional 2-center, 2-electron covalent bonds. The bonding between the boron atoms and the bridging hydrogen atoms is, however, different from that in molecules such as hydrocarbons. Having used two electrons in bonding to the terminal hydrogen atoms, each boron has one valence electron remaining for additional bonding. The bridging hydrogen atoms provide one electron each. Thus the BH ring is held together by four electrons, an example of 3-center 2-electron bonding. This type of bond is sometimes called a 'banana bond'. The lengths of the B-H bonds and the B-H bonds are 1.33 and 1.19 Å respectively, and this difference in the lengths of these bonds reflects the difference in their strengths, the B-H bonds being relatively weaker. The weakness of the B-H vs B-H bonds is indicated by their vibrational signatures in the infrared spectrum, being ~2100 and 2500 cm, respectively. The structure is isoelectronic with CH, which would arise from the diprotonation of the planar molecule ethene. Diborane is one of many compounds with such unusual bonding.
Resonance-assisted hydrogen bond
It is caused 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.
Hydrogen bond
Symmetric hydrogen bonds have been observed recently spectroscopically in formic acid at high pressure (>GPa). Each hydrogen atom forms a partial covalent bond with two atoms rather than one. Symmetric hydrogen bonds have been postulated in ice at high pressure (Ice X). Low-barrier hydrogen bonds form when the distance between two heteroatoms is very small.
Carbon–oxygen bond
Carbon and oxygen form terminal double bonds in functional groups collectively known as carbonyl compounds to which belong such compounds as ketones, esters, carboxylic acids and many more. Internal C=O bonds are found in positively charged oxonium ions. In furans, the oxygen atom contributes to pi-electron delocalization via its filled p-orbital and hence furans are aromatic. Bond lengths of C=O bonds are around 123 pm in carbonyl compounds. The C=O bond length in carbon dioxide is 116 pm. The C=O bonds in acyl halides have partial triple bond character and are subsequently very short: 117 pm. Compounds with formal C–O triple bonds do not exist except for carbon monoxide, which has a very short, strong bond (112.8 pm). Such triple bonds have a very high bond energy, even higher than N–N triple bonds./ref> Oxygen can also be trivalent, for example in triethyloxonium tetrafluoroborate.
TEMPO
Regardless of the reasons for the stability of the radical, the O–H bond in the hydrogenated derivative (the hydroxylamine) TEMPO–H is weak. With an O–H bond dissociation energy of about 70 kcal/mol, this bond is about 30% weaker than a typical O–H bond.
Double bond
With 133 pm, the ethylene C=C bond length is shorter than the C−C length in ethane with 154 pm. The double bond is also stronger, 636 kJ mol versus 368 kJ mol but not twice as much as the pi-bond is weaker than the sigma bond due to less effective pi-overlap.
Staebler–Wronski effect
Historically, the most favored model has been the hydrogen bond switching model. It proposes that an electron-hole pair formed by the incident light may recombine near a weak Si–Si bond, releasing energy sufficient to break the bond. A neighbouring H atom then forms a new bond with one of the Si atoms, leaving a dangling bond. These dangling bonds can trap electron-hole pairs, thus reducing the current that can pass through. However, new experimental evidence is casting doubt on this model. More recently, the "H collision model" proposed that two spatially separated recombination events cause emission of mobile hydrogen from Si–H bonds to form two dangling bonds, with a metastable paired H state binding the hydrogen atoms at a distant site.
Halogen bond
However, it was not until the mid-1990s, that the nature and applications of the halogen bond began to be intensively studied. Systematic and extensive work by Legon and coworkers, who studied a wide variety of halogen-bonded adducts formed in the gas phase via microwave spectroscopy, drew attention to the similarities between halogen-bonding and better-known hydrogen-bonding interactions. Computational calculations by Politzer and Murray were particularly significant as they revealed that the high directionality of the halogen bond is the result of an anisotropic distribution of electron density around the halogen nucleus and paved the way to the definition of the “σ-hole”.
Drug action
There is a small but significant attraction between hydrogen atoms and polar functional groups (e.g. the hydroxyl [-OH] group). These so-called hydrogen bonds only act over short distances, and are dependent on the correct alignment between functional groups.
Halogen bond
Aldehyde-based and carboxylate inhibitors are effective but toxic because the functional activity of aldehyde reductase is impaired. Carboxylate and aldehyde inhibitors were shown to hydrogen bond with Trp 111, Tyr 48, and His 110. The “specificity pocket,” created as a result of inhibitor binding, consists of Leu 300, Ala 299, Phe 122, Thr 113, and Trp 111. For inhibitors to be effective, the key residues of interaction were identified to be Thr 113 and Trp 111. IDD 594 was designed such that the halogen would provide selectivity and be potent. Upon binding, this compound induces a conformational change that causes halogen bonding to occur between the oxygen of the Thr and the bromine of the inhibitor. The bond distance was measured to be 2.973(4) Å. It is this O−Br halogen bond that contributes to the large potency of this inhibitor for human aldose reductase rather than aldehyde reductase.
Pectin
In high-ester pectins at soluble solids content above 60% and a pH-value between 2.8 and 3.6, hydrogen bonds and hydrophobic interactions bind the individual pectin chains together. These bonds form as water is bound by sugar and forces pectin strands to stick together. These form a 3-dimensional molecular net that creates the macromolecular gel. The gelling-mechanism is called a low-water-activity gel or sugar-acid-pectin gel.
Hydrochloric acid
As a strong acid, hydrogen chloride has a large "K". Theoretical attempts to assign the p"K" of hydrogen chloride have been made, with the most recent estimate being −5.9. However, it is important to distinguish between hydrogen chloride gas and hydrochloric acid. Due to the leveling effect, except when highly concentrated and behavior deviates from ideality, "hydrochloric acid" (aqueous HCl) is only as acidic as the strongest proton donor available in water, the aquated proton (popularly known as "hydronium ion"). When chloride salts such as NaCl are added to aqueous HCl, they have only a minor effect on pH, indicating that Cl is a very weak conjugate base and that HCl is fully dissociated in aqueous solution. Dilute solutions of HCl have a pH close to that predicted by assuming full dissociation into hydrated H and Cl.
Hydrogen bond
A protein backbone hydrogen bond incompletely shielded from water attack is a dehydron. Dehydrons promote the removal of water through proteins or ligand binding. The exogenous dehydration enhances the electrostatic interaction between the amide and carbonyl groups by de-shielding their partial charges. Furthermore, the dehydration stabilizes the hydrogen bond by destabilizing the nonbonded state consisting of dehydrated isolated charges.
Base pair
Hydrogen bonding is the chemical interaction that underlies the base-pairing rules described above. Appropriate geometrical correspondence of hydrogen bond donors and acceptors allows only the "right" pairs to form stably. DNA with high GC-content is more stable than DNA with low GC-content. But, contrary to popular belief, the hydrogen bonds do not stabilize the DNA significantly; stabilization is mainly due to stacking interactions.
Direct bonding
Between room temperature and 110 °C the interface energy remains low, water molecules diffuse at the bond interface, leading to a rearrangement, causing more hydrogen-bonds. At temperatures from 110 °C to 150 °C silanol groups polymerize to siloxane and water, but also a slow fracture takes place. This reaction equates a thermo dynamical equilibrium and a higher density of silanol groups results in a higher number of siloxane and an increasing bond strength.
Helium compounds
The HeHN molecule is linear. The He-H bondlength is 1.72 Å. It has an infrared band, due to B-H stretching, with a base at 3158.42 cm. The binding energy is 378 cm in the 000 vibrational state, and 431 cm in the 100 vibrational state. HeHN is also known. One helium atom is linked to a hydrogen, and the other is less tightly bound.
Human bonding
In 1962, while a freshman history major at Harvard, Mark Granovetter became enamored of the concepts underlying the classic chemistry lecture in which "weak" hydrogen bonds hold huge numbers of water molecules together, which themselves are held together by "strong" covalent bonds. This model was the stimulus behind his famous 1973 paper "The Strength of Weak Ties", which is now considered a classic paper in sociology.
Hydrogen bond
The hydrogen bond is often described as an electrostatic dipole-dipole interaction. However, it also has some features of covalent bonding: it is directional and strong, produces interatomic distances shorter than the sum of the van der Waals radii, and usually involves a limited number of interaction partners, which can be interpreted as a type of valence. These covalent features are more substantial when acceptors bind hydrogens from more electronegative donors.
Iron hydride
Iron and iron-based alloys can form solid solutions with hydrogen, which under extreme pressure may reach stoichiometric proportions, remaining stable even at high temperatures and that is reported to survive for a while under ambient pressure, at temperatures below 150K.
Complexes displaying iron–hydrogen bonds include, for example:
Organoiodine compound
The C–I bond is the weakest of the carbon–halogen bonds. These bond strengths correlate with the electronegativity of the halogen, decreasing in the order F > Cl > Br > I. This periodic order also follows the atomic radius of halogens and the length of the carbon-halogen bond. For example, in the molecules represented by CHX, where X is a halide, the carbon-X bonds have strengths, or bond dissociation energies, of 115, 83.7, 72.1, and 57.6 kcal/mol for X = fluoride, chloride, bromide, and iodide, respectively. Of the halides, iodide usually is the best leaving group. Because of the weakness of the C-I bond, samples of organoiodine compounds are often yellow due to an impurity of I.
Low-barrier hydrogen bond
A Low-barrier hydrogen bond (LBHB) is a special type of hydrogen bond. LBHBs can occur when the pKa of the two heteroatoms are closely matched, which allows the hydrogen to be more equally shared between them. This hydrogen-sharing causes the formation of especially short, strong hydrogen bonds.
Hydride
Bonds between hydrogen and other elements range from highly to somewhat covalent. Some hydrides, e.g. boron hydrides, do not conform to classical electron-counting rules, and the bonding is described in terms of multi-centered bonds, whereas the interstitial hydrides often involve metallic bonding. Hydrides can be discrete molecules, oligomers or polymers, ionic solids, chemisorbed monolayers, bulk metals (interstitial), and other materials. While hydrides traditionally react as Lewis bases or reducing agents, some metal hydrides behave as hydrogen-atom donors and act as acids.
Free hydride anions exist only under extreme conditions and are not invoked for homogeneous solution. Instead, many compounds have hydrogen centres with hydridic character.
Chemical bond
In nonpolar covalent bonds, the electronegativity difference between the bonded atoms is small, typically 0 to 0.3. Bonds within most organic compounds are described as covalent. The figure shows methane (CH), in which each hydrogen forms a covalent bond with the carbon. See sigma bonds and pi bonds for LCAO-description of such bonding.
Hydrogen bond
Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule). Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, and environment, the energy of a hydrogen bond can vary between 1 and 40 kcal/mol. This makes them somewhat stronger than a van der Waals interaction, and weaker than fully covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins.
Peptide bond
A peptide bond can be broken by hydrolysis (the addition of water). In the presence of water they will break down and release 8–16 kilojoule/mol (2–4 kcal/mol) of Gibbs energy. This process is extremely slow, with the half life at 25 °C of between 350 and 600 years per bond.
Hydrogen bond
Liquid water's high boiling point is due to the high number of hydrogen bonds each molecule can form, relative to its low molecular mass. Owing to the difficulty of breaking these bonds, water has a very high boiling point, melting point, and viscosity compared to otherwise similar liquids not conjoined by hydrogen bonds. Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is up to four.
Hydrogen bond
A ubiquitous example of a hydrogen bond is found between water molecules. In a discrete water molecule, there are two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them that is to say oxygen-hydrogen bonding; the simplest case, when only two molecules are present, is called the water dimer and is often used as a model system. When more molecules are present, as is the case with liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with a hydrogen on another water molecule. This can repeat such that every water molecule is H-bonded with up to four other molecules, as shown in the figure (two through its two lone pairs, and two through its two hydrogen atoms). Hydrogen bonding strongly affects the crystal structure of ice, helping to create an open hexagonal lattice. The density of ice is less than the density of water at the same temperature; thus, the solid phase of water floats on the liquid, unlike most other substances.
Bond cleavage
In homolytic cleavage, or homolysis, the two electrons in a cleaved covalent bond are divided equally between the products. This process is also known as "homolytic fission" or "radical fission". The bond-dissociation energy of a bond is the amount of energy required to cleave the bond homolytically. This enthalpy change is one measure of bond strength.
Charge-shift bond
The valence bond view of chemical bonding that owes much to the work of Pauling is familiar to many, if not all, chemists. The basis of Pauling's description of the chemical bond is that an electron pair bond involves the mixing, resonance, of one covalent and two ionic structures. In bonds between two atoms of the same element, homonuclear bonds, Pauling assumed that the ionic structures make no appreciable contribution to the overall bonding. This assumption followed on from published calculations for the hydrogen molecule in 1933 by Weinbaum and by James and Coolidge. that showed that the contribution of ionic forms amounted to only a small percentage of the H−H bond energy. For heteronuclear bonds, A−X, Pauling estimated the covalent contribution to the bond dissociation energy as being the mean of the bond dissociation energies of homonuclear A−A and X−X bonds. The difference between the mean and the observed bond energy was assumed to be due to the ionic contribution. The calculation for HCl is shown below.
Agostic interaction
Agostic interactions are best demonstrated by crystallography. Neutron diffraction data has shown that C−H and M┄H bond distances are 5-20% longer than expected for isolated metal hydride and hydrocarbons. The distance between the metal and the hydrogen is typically 1.8–2.3 Å, and the M┄H−C angle falls in the range 90°–140°. The presence of a H NMR signal that is shifted upfield from that of a normal aryl or alkane, often to the region normally assigned to hydride ligands. The coupling constant "J" is typically lowered to 70–100 Hz versus the 125 Hz expected for a normal sp carbon–hydrogen bond.
On the basis of experimental and computational studies, the stabilization arising from an agostic interaction is estimated to be 10–15 kcal/mol. Recent calculations using compliance constants point to a weaker stabilisation (<10 kcal/mol). Thus, agostic interactions are stronger than most hydrogen bonds. Agostic bonds sometimes play a role in catalysis by increasing 'rigidity' in transition states. For instance, in Ziegler–Natta catalysis the highly electrophilic metal center has agostic interactions with the growing polymer chain. This increased rigidity influences the stereoselectivity of the polymerization process.
Water
The hydrogen bonds are also the reason why the melting and boiling points of water are much higher than those of other analogous compounds like hydrogen sulfide (). They also explain its exceptionally high specific heat capacity (about 4.2 J/g/K), heat of fusion (about 333 J/g), heat of vaporization (), and thermal conductivity (between 0.561 and 0.679 W/m/K). These properties make water more effective at moderating Earth's climate, by storing heat and transporting it between the oceans and the atmosphere. The hydrogen bonds of water are of moderate strength, around 23 kJ/mol (compared to a covalent O-H bond at 492 kJ/mol). Of this, it is estimated that 90% of the hydrogen bond is attributable to electrostatics, while the remaining 10% reflects partial covalent character.
Bond valence method
Structures containing covalent bonds can be treated using the ionic model providing they satisfy the topological conditions given above, but a special situation applies to hydrocarbons which allows the bond valence model to be reduced to the traditional bond model of organic chemistry. If an atom has a valence, V, that is equal to its coordination number, N, its bonding strength according to Eq. 3 is exactly 1.0 vu (valence units), a condition that greatly simplifies the model. This condition is obeyed by carbon, hydrogen and silicon. Since these atoms all have bonding strengths of 1.0 vu the bonds between them are all predicted to have integral valences with carbon forming four single bonds and hydrogen one. Under these conditions, the bonds are all single bonds (or multiples of single bonds). Compounds can be constructed by linking carbon and hydrogen atoms with bonds that are all exactly equivalent. Under certain conditions, nitrogen can form three bonds and oxygen two, but since nitrogen and oxygen typically also form hydrogen bonds, the resulting N-H and O-H bonds have valences less than 1.0 vu, leading through the application of Eq. 1, to the C-C and C-H bonds having valences that differ from 1.0 vu. Nevertheless, the simple bonding rules of organic chemistry are still good approximations, though the rules of the bond valence model are better.
Base (chemistry)
A strong base is a basic chemical compound that can remove a proton (H) from (or "deprotonate") a molecule of even a very weak acid (such as water) in an acid-base reaction. Common examples of strong bases include hydroxides of alkali metals and alkaline earth metals, like NaOH and , respectively. Due to their low solubility, some bases, such as alkaline earth hydroxides, can be used when the solubility factor is not taken into account. One advantage of this low solubility is that "many antacids were suspensions of metal hydroxides such as aluminum hydroxide and magnesium hydroxide." These compounds have low solubility and have the ability to stop an increase in the concentration of the hydroxide ion, preventing the harm of the tissues in the mouth, oesophagus, and stomach. As the reaction continues and the salts dissolve, the stomach acid reacts with the hydroxide produced by the suspensions. Strong bases hydrolyze in water almost completely, resulting in the leveling effect." In this process, the water molecule combines with a strong base, due to the water's amphoteric ability; and, a hydroxide ion is released. Very strong bases can even deprotonate very weakly acidic C–H groups in the absence of water. Here is a list of several strong bases:The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals).
Ice Ih
The hydrogen atoms in the crystal lattice lie very nearly along the hydrogen bonds, and in such a way that each water molecule is preserved. This means that each oxygen atom in the lattice has two hydrogens adjacent to it, at about 101 pm along the 275 pm length of the bond. The crystal lattice allows a substantial amount of disorder in the positions of the hydrogen atoms frozen into the structure as it cools to absolute zero. As a result, the crystal structure contains some residual entropy inherent to the lattice and determined by the number of possible configurations of hydrogen positions that can be formed while still maintaining the requirement for each oxygen atom to have only two hydrogens in closest proximity, and each H-bond joining two oxygen atoms having only one hydrogen atom. This residual entropy "S" is equal to 3.5 J mol K.
Alkene
Like a single covalent bond, double bonds can be described in terms of overlapping atomic orbitals, except that, unlike a single bond (which consists of a single sigma bond), a carbon–carbon double bond consists of one sigma bond and one pi bond. This double bond is stronger than a single covalent bond (611 kJ/mol for C=C vs. 347 kJ/mol for C–C) and also shorter, with an average bond length of 1.33 ångströms (133 pm).
Bromine
At room temperature, hydrogen bromide is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative bromine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen bromide at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised. Aqueous hydrogen bromide is known as hydrobromic acid, which is a strong acid (p"K" = −9) because the hydrogen bonds to bromine are too weak to inhibit dissociation. The HBr/HO system also involves many hydrates HBr·"n"HO for "n" = 1, 2, 3, 4, and 6, which are essentially salts of bromine anions and hydronium cations. Hydrobromic acid forms an azeotrope with boiling point 124.3 °C at 47.63 g HBr per 100 g solution; thus hydrobromic acid cannot be concentrated beyond this point by distillation.
Organosilicon
Si–O bonds are much stronger (809 kJ/mol compared to 538 kJ/mol) than a typical C–O single bond. The favorable formation of Si–O bonds drives many organic reactions such as the Brook rearrangement and Peterson olefination. Compared to the strong Si–O bond, the Si–F bond is even stronger.
Active site
Hydrogen bond: A hydrogen bond is a specific type of dipole-dipole interaction between a partially positive hydrogen atom and a partially negative electron donor that contain a pair of electrons such as oxygen, sulfur and nitrogen. The strength of hydrogen bond depends on the chemical nature and geometric arrangement of each group.
Dispersive adhesion
Sometimes grouped into the chemical mechanism of adhesion, hydrogen bonding can increase adhesive strength by the dispersive mechanism. Hydrogen bonding occurs between molecules with a hydrogen atom attached to a small, electronegative atom such as fluorine, oxygen or nitrogen. This bond is naturally polar, with the hydrogen atom gaining a slight positive charge and the other atom becoming slightly negative. Two molecules, or even two functional groups on one large molecule, may then be attracted to each other via Keesom forces.
Hydrogen bond
A symmetric hydrogen bond is a special type of hydrogen bond in which the proton is spaced exactly halfway between two identical atoms. The strength of the bond to each of those atoms is equal. It is an example of a three-center four-electron bond. This type of bond is much stronger than a "normal" hydrogen bond. The effective bond order is 0.5, so its strength is comparable to a covalent bond. It is seen in ice at high pressure, and also in the solid phase of many anhydrous acids such as hydrofluoric acid and formic acid at high pressure. It is also seen in the bifluoride ion [F−H−F].
Trihydrogen cation
The three hydrogen atoms in the molecule form an equilateral triangle, with a bond length of 0.90 Å on each side. The bonding among the atoms is a three-centre two-electron bond, a delocalized resonance hybrid type of structure. The strength of the bond has been calculated to be around 4.5 eV (104 kcal/mol).
Pnictogen hydride
All straight-chain saturated hydrogen pnictides follow the formula XH.
Ammonia is produced industrially on the largest scale among all compounds. Like water, hydrogen bonding results in a high melting and boiling point compared to the other pnictogen hydrides, although 26% is lost on melting, another 7% as the liquid is heated to boiling, and the remaining 67% upon boiling. Other effects of hydrogen bonding are a high dielectric constant as well as low values of density, viscosity, and electrical conductivity. Like water, it is an excellent and often-used ionising solvent. Over twenty other hydrides of nitrogen are known, the most important being hydrazine (NH) and hydrogen azide (HN). Hydrazine has physical properties that are remarkably similar to those of water: its melting and boiling points are 2.0 °C and 113.5 °C, the density of the solid at −5 °C is 1.146 g/cm, while that of the liquid at 25 °C is 1.00 g/cm. The azanes are a series which include ammonia, hydrazine and triazane.
Alkyne
The triple bond is very strong with a bond strength of 839 kJ/mol. The sigma bond contributes 369 kJ/mol, the first pi bond contributes 268 kJ/mol and the second pi-bond of 202 kJ/mol bond strength. Bonding usually discussed in the context of molecular orbital theory, which recognizes the triple bond as arising from overlap of s and p orbitals. In the language of valence bond theory, the carbon atoms in an alkyne bond are sp hybridized: they each have two unhybridized p orbitals and two sp hybrid orbitals. Overlap of an sp orbital from each atom forms one sp–sp sigma bond. Each p orbital on one atom overlaps one on the other atom, forming two pi bonds, giving a total of three bonds. The remaining sp orbital on each atom can form a sigma bond to another atom, for example to hydrogen atoms in the parent acetylene. The two sp orbitals project on opposite sides of the carbon atom.
Properties of water
A single water molecule can participate in a maximum of four hydrogen bonds because it can accept two bonds using the lone pairs on oxygen and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia and methanol can also form hydrogen bonds. However, they do not show anomalous thermodynamic, kinetic or structural properties like those observed in water because none of them can form four hydrogen bonds: either they cannot donate or accept hydrogen atoms, or there are steric effects in bulky residues. In water, intermolecular tetrahedral structures form due to the four hydrogen bonds, thereby forming an open structure and a three-dimensional bonding network, resulting in the anomalous decrease in density when cooled below 4 °C. This repeated, constantly reorganizing unit defines a three-dimensional network extending throughout the liquid. This view is based upon neutron scattering studies and computer simulations, and it makes sense in the light of the unambiguously tetrahedral arrangement of water molecules in ice structures.
Protein secondary structure
Secondary structure is defined by hydrogen bonding, so the exact definition of a hydrogen bond is critical. The standard hydrogen-bond definition for secondary structure is that of DSSP, which is a purely electrostatic model. It assigns charges of ±"q" ≈ 0.42"e" to the carbonyl carbon and oxygen, respectively, and charges of ±"q" ≈ 0.20"e" to the amide hydrogen and nitrogen, respectively. The electrostatic energy is
Low-barrier hydrogen bond
Standard hydrogen bonds are longer (e.g. 2.8 Å for an O···O h-bond), and the hydrogen ion clearly belongs to one of the heteroatoms. When pKa of the heteroatoms is closely matched, a LBHB becomes possible at a shorter distance (~2.55 Å). When the distance further decreases (< 2.29 Å) the bond is characterized as a single-well or short-strong hydrogen bond.
Hydrogen bond
For example, hydrogen fluoride—which has three lone pairs on the F atom but only one H atom—can form only two bonds; (ammonia has the opposite problem: three hydrogen atoms but only one lone pair).Hydrogen bonding plays an important role in determining the three-dimensional structures and the properties adopted by many synthetic and natural proteins.
Alkali metal
The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher than that of the alkali metals. As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative hydride ion, and is very occasionally considered to be a halogen on that basis. (The alkali metals can also form negative ions, known as alkalides, but these are little more than laboratory curiosities, being unstable.) An argument against this placement is that formation of hydride from hydrogen is endothermic, unlike the exothermic formation of halides from halogens. The radius of the H anion also does not fit the trend of increasing size going down the halogens: indeed, H is very diffuse because its single proton cannot easily control both electrons. It was expected for some time that liquid hydrogen would show metallic properties; while this has been shown to not be the case, under extremely high pressures, such as those found at the cores of Jupiter and Saturn, hydrogen does become metallic and behaves like an alkali metal; in this phase, it is known as metallic hydrogen. The electrical resistivity of liquid metallic hydrogen at 3000 K is approximately equal to that of liquid rubidium and caesium at 2000 K at the respective pressures when they undergo a nonmetal-to-metal transition.
Silyl hydride
Silicon hydrides are chemical compounds which contain a silicon–hydrogen bond. The silicon-to-hydrogen bond is longer than the C–H bond (148 compared to 105 pm) and weaker (299 compared to 338 kJ/mol). Hydrogen is more electronegative than silicon (hence the naming convention of silyl hydrides), which results in the polarization of the Si-H bond to be the reverse of that for the C-H bond. Generally silyl hydrides are colourless with physical properties (solubility, volatility) comparable to hydrocarbons. They can be pyrophoric, reflecting the great driving force for replacing Si-H bonds with Si-O bonds.
Organosilicon
The silicon to hydrogen bond is longer than the C–H bond (148 compared to 105 pm) and weaker (299 compared to 338 kJ/mol). Hydrogen is more electronegative than silicon hence the naming convention of silyl "hydrides". Commonly the presence of the hydride is not mentioned in the name of the compound. Triethylsilane has the formula EtSiH. Phenylsilane is PhSiH. The parent compound SiH is called silane.
Chemical bonding of H2O
The Lewis structure of describes the bonds as two sigma bonds between the central oxygen atom and the two peripheral hydrogen atoms with oxygen having two lone pairs of electrons. Valence bond theory suggests that is sp3 hybridized in which the 2s atomic orbital and the three 2p orbitals of oxygen are hybridized to form four new hybridized orbitals which then participate in bonding by overlapping with the hydrogen 1s orbitals. As such, the predicted shape and bond angle of sp3 hybridization is tetrahedral and 109.5°. This is in open agreement with the true bond angle of 104.45°. The difference between the predicted bond angle and the measured bond angle is traditionally explained by the electron repulsion of the two lone pairs occupying two sp3 hybridized orbitals. While valence bond theory is suitable for predicting the geometry and bond angle of , its prediction of electronic states does not agree with the experimentally measured reality. In the valence bond model, the two sigma bonds are of identical energy and so are the two lone pairs since they both resides in the same bonding and nonbonding orbitals, thus corresponding to two energy levels in the photoelectronic spectrum. In other words, if water was formed from two identical O-H bonds and two identifcal sp3 lone pairs on the oxygen atom as predicted by valence bond theory, then its photoelectron spectrum (PES) would have two (degenerate) peaks and energy, one for the two O-H bonds and the other for the two sp lone pairs. However, the photoelectronic spectrum of reveals four different energy levels that correspond to the ionization energies of the two bonding and two nonbonding pairs of elections at 12.6eV, 14.7eV, 18.5eV, and 32.2eV. This suggest that neither the two O-H bonds nor the two sp3 lone pairs are degenerate in energy.
Hydrogen embrittlement
Although hydrogen atoms embrittle a variety of substances, including steel, aluminium (at high temperatures only), and titanium, however these metals are still affected in high concentrations, hydrogen embrittlement of high-strength steel is of the most importance. Austempered iron is also susceptible, though austempered steel (and possibly other austempered metals) display increased resistance to hydrogen embrittlement. Steel with an ultimate tensile strength of less than 1000 MPa (~145,000 psi) or hardness of less than 23 HRC is not generally considered susceptible to hydrogen embrittlement. In tensile tests carried out on several structural metals under high-pressure molecular hydrogen environment, it has been shown that austenitic stainless steels, aluminium (including alloys), copper (including alloys, e.g. beryllium copper) are not susceptible to hydrogen embrittlement along with a few other metals. As an example of severe hydrogen embrittlement, the elongation at failure of 17-4PH precipitation hardened stainless steel was measured to drop from 17% to only 1.7% when smooth specimens were exposed to high-pressure hydrogen.
Organic peroxide
The O−O bond length in peroxides is about 1.45 Å, and the R−O−O angles (R = H, C) are about 110° (water-like). Characteristically, the C−O−O−R (R = H, C) dihedral angles are about 120°. The O−O bond is relatively weak, with a bond dissociation energy of , less than half the strengths of C−C, C−H, and C−O bonds.
Wafer bond characterization
The bond strength can be evaluated using double cantilever beam or chevron respectively micro-chevron tests. Other pull tests as well as burst, direct shear tests or bend tests enable the determination of the bond strength. The packaging hermeticity is characterized using membrane, He-leak, resonator/pressure tests.
Metal carbonyl hydride
The neutral metal carbonyl hydrides are often volatile and can be quite acidic. The hydrogen atom is directly bounded to the metal. The metal-hydrogen bond length is for cobalt 114 pm, the metal-carbon bond length is for axial ligands 176  and 182  for the equatorial ligands.
Halogen bond
To test the notion of polarizability involvement in the strength of halogen bonding, bromopentafluorbenzene was used as a Lewis base. Consequently, verification of halogen bond complex formation wasn’t obtained. This finding provides more support for the dependence of halogen bonding on atomic polarizability. Utilizing similar donor-acceptor frameworks, the authors demonstrated that halogen bonding strength in the liquid crystalline state is comparable to the hydrogen-bonded mesogens.
9,10-Dihydroanthracene
The bond dissociation energy for the 9- and 10- carbon–hydrogen bonds are estimated at 78 kcal mol. Thus these bonds are about 20% weaker than typical C–H bonds.
Bond cleavage
The singlet excitation energy of a sigma bond is the energy required for heterolytic dissociation, but the actual singlet excitation energy may be lower than the bond dissociation energy of heterolysis as a result of the Coulombic attraction between the two ion fragments. The singlet excitation energy of a silicon–silicon sigma bond is lower than the carbon–carbon sigma bond, even though their bond strengths are 80kJ/mol and 70kJ/mol respectively, because silicon has higher electron affinity and lower ionization potential than carbon.
Molecular solid
A hydrogen bond is a specific dipole where a hydrogen atom has a partial positive charge (δ+) to due a neighboring electronegative atom or functional group. Hydrogen bonds are amongst the strong intermolecular interactions know other than ion-dipole interactions. For intermolecular hydrogen bonds the δ+ hydrogen interacts with a δ- on an adjacent molecule. Examples of molecular solids that hydrogen bond are water, amino acids, and acetic acid. For acetic acid, the hydrogen (δ+) on the alcohol moiety of the carboxylic acid hydrogen bonds with other the carbonyl moiety (δ-) of the carboxylic on the adjacent molecule. This hydrogen bond leads a string of acetic acid molecules hydrogen bonding to minimize free energy. These strings of acetic acid molecules then stack together to build solids.
Hydrogen bond
Quantum chemical calculations of the relevant interresidue potential constants (compliance constants) revealed 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.
Hydrogen bond
Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have much weaker hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.
Nitrogen
There are some theoretical indications that other nitrogen oligomers and polymers may be possible. If they could be synthesised, they may have potential applications as materials with a very high energy density, that could be used as powerful propellants or explosives. This is because they should all decompose to dinitrogen, whose N≡N triple bond (bond energy 946 kJ⋅mol) is much stronger than those of the N=N double bond (418 kJ⋅mol) or the N–N single bond (160 kJ⋅mol): indeed, the triple bond has more than thrice the energy of the single bond. (The opposite is true for the heavier pnictogens, which prefer polyatomic allotropes.) A great disadvantage is that most neutral polynitrogens are not expected to have a large barrier towards decomposition, and that the few exceptions would be even more challenging to synthesise than the long-sought but still unknown tetrahedrane. This stands in contrast to the well-characterised cationic and anionic polynitrogens azide (), pentazenium (), and pentazolide (cyclic aromatic ). Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced in a diamond anvil cell, nitrogen polymerises into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds, resulting in its nickname "nitrogen diamond".
At atmospheric pressure, molecular nitrogen condenses (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha phase). Liquid nitrogen, a colourless fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen. Solid nitrogen has many crystalline modifications. It forms a significant dynamic surface coverage on Pluto and outer moons of the Solar System such as Triton. Even at the low temperatures of solid nitrogen it is fairly volatile and can sublime to form an atmosphere, or condense back into nitrogen frost. It is very weak and flows in the form of glaciers and on Triton geysers of nitrogen gas come from the polar ice cap region.
Hydrogen bond
Hydrogen bonding is of continuing theoretical interest. 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.
Mercury(II) hydride
The HgH molecule is linear and symmetric in the form H-Hg-H. The bond length is 1.646543 Å. The antisymmetric stretching frequency, ν of the bond is 1912.8 cm, 57.34473 THz for isotopes Hg and H. The energy needed to break the Hg-H bond in HgH is 70 kcal/mol. The second bond in the resulting HgH is much weaker only needing 8.6 kcal/mol to break. Reacting two hydrogen atoms releases 103.3 kcal/mol, and so HgH formation from hydrogen molecules and Hg gas is endothermic at 24.2 kcal/mol.