What is the strength of 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:
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.
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.
N-heterocyclic silylene
Atoms in Molecules bond energy at critical point calculations show that the N to Si bonds in an NHSi are significantly ionic, unlike the highly covalent N-C bond in an NHC. Furthermore, the charge density and the Laplacian (∇) at the bond critical points between the silylene and carbene are dissimilar. Silylenes have much lower charge density between Si and N, with a positive Laplacian from Si to N indicating a more ionic interaction, while the Laplacian is negative going from carbon to nitrogen in an NHC indicating a more covalent bond. Upon coordination to a Group 11 metal the covalent character in the Si or C to metal bond is small, but not insignificant, again with a higher covalent character in the carbenic system. Similarly, silicon has a higher Laplacian between it and the group 11 metals than does carbon, indicating a less covalent metal interaction in silyenes as compared to carbenes.
Glass frit bonding
The bonding strength is mainly dependent on the density, the spreading area of the glass frit layer and the surface layer of the bonding interface. It is high enough, around 20 MPa, for most applications and comparable to those achieved with anodic bonding. The hermeticity ensures the correct function and a sufficient reliability of the bond and therefore the product. Further, the bonding yield of glass frit bonded wafers is very high, normally > 90 %.
Alkene
Each carbon of the double bond uses its three sp hybrid orbitals to form sigma bonds to three atoms (the other carbon and two hydrogen atoms). The unhybridized 2p atomic orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid orbitals, combine to form the pi bond. This bond lies outside the main C–C axis, with half of the bond on one side of the molecule and half on the other. With a strength of 65 kcal/mol, the pi bond is significantly weaker than the sigma bond.
Hydrogen bond
In the secondary structure of proteins, hydrogen bonds form between the backbone oxygens and amide hydrogens. When the spacing of the amino acid residues participating in a hydrogen bond occurs regularly between positions "i" and "i" + 4, an alpha helix is formed. When the spacing is less, between positions "i" and "i" + 3, then a 3 helix is formed. When two strands are joined by hydrogen bonds involving alternating residues on each participating strand, a beta sheet is formed. Hydrogen bonds also play a part in forming the tertiary structure of protein through interaction of R-groups. (See also protein folding).
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.
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).
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.
Hydrogen bond
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 indicases, using the electron density of the system.
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.
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.
Carbon–hydrogen bond
The C−H bond in general is very strong, so it is relatively unreactive. In several compound classes, collectively called carbon acids, the C−H bond can be sufficiently acidic for proton removal. Unactivated C−H bonds are found in alkanes and are not adjacent to a heteroatom (O, N, Si, etc.). Such bonds usually only participate in radical substitution. Many enzymes are known, however, to affect these reactions.
Chemical bond
Strong chemical bonds are the "intramolecular" forces which hold atoms together in molecules. A strong chemical bond is formed from the transfer or sharing of electrons between atomic centers and relies on the electrostatic attraction between the protons in nuclei and the electrons in the orbitals.
1.1.1-Propellane
[1.1.1]Propellane is a highly strained molecule. The bonds of the two central carbon atoms have an inverted tetrahedral geometry, and the length of the central bond is 160 pm. The strength of that bond is disputed; estimates vary from 59–65 kcal/mol to no strength at all. The energy of the biradical state (with no central bond at all) is calculated to be 80 kcal/mol higher. The compound is highly unstable, and at 114 °C it will spontaneously isomerize to 3-methylidenecyclobutene with a half-life of 5 minutes. Its strain energy is estimated to be 102 kcal/mol (427 kJ/mol). Surprisingly, [1.1.1]propellane is persistent at room temperature and is somewhat less susceptible to thermal decomposition than the less strained (90 kcal/mol) [2.2.2]propellane system, which has an estimated half-life of only about 1 h at 25 °C.
Carbon–hydrogen bond
Although the C−H bond is one of the strongest, it varies over 30% in magnitude for fairly stable organic compounds, even in the absence of heteroatoms.
The C−H unit is formally named carbogen.
Hydrogen bond
A single hydrogen atom can participate in two hydrogen bonds, rather than one. This type of bonding is called "bifurcated" (split in two or "two-forked"). It can exist, for instance, in complex natural or synthetic organic molecules. It has been suggested that a bifurcated hydrogen atom is an essential step in water reorientation.
Acceptor-type hydrogen bonds (terminating on an oxygen's lone pairs) are more likely to form bifurcation (it is called overcoordinated oxygen, OCO) than are donor-type hydrogen bonds, beginning on the same oxygen's hydrogens.
Bond-dissociation energy
The bond dissociation energy is an enthalpy change of a particular chemical process, namely homolytic bond cleavage, and "bond strength" as measured by the BDE should not be regarded as an intrinsic property of a particular bond type but rather as an energy change that depends on chemical context. For instance, Blanksby and Ellison cites the example of ketene (HC=CO), which has a C=C bond dissociation energy of 79 kcal/mol, while ethylene (HC=CH) has a bond dissociation energy of 174 kcal/mol. This vast difference is accounted for by the thermodynamic stability of carbon monoxide (CO), formed upon C=C bond cleavage of ketene. The difference in availability of spin states upon fragmentation further complicates the use of BDE as a measure of bond strength for head-to-head comparisons, and force constants have been suggested as an alternative.
Iron–hydrogen alloy
The hydrogen in typical iron hydrides may contribute up to 13 ppm in its weight. Varying the amount of hydrogen, as well as controlling its chemical and physical makeup in the final iron hydride (either as a solute element, or as a precipitated phase), hastens the movement of those dislocations that make pure iron ductile, and thus controls and undermines its qualities. Varying the other alloying elements and controlling their chemical and physical makeup also controls, but enhances its qualities. These qualities include such things as the hardness, quenching behaviour, need for annealing, tempering behaviour, yield strength, and tensile strength of the resulting iron-hydrogen alloy. The retension of iron hydride's strength compared to pure iron is possible only by maintaining iron's ductility.
Halogen bond
On the bio-molecular level, halogen bonding is important for substrate specificity, binding and molecular folding. In the case of protein-ligand interactions, the most common charge-transfer bonds with polarizable halogens involve backbone carbonyls and/or hydroxyl and carboxylate groups of amino acid residues. Typically in DNA and protein-ligand complexes, the bond distance between Lewis base donor atoms (e.g. O, S, N) and Lewis acid (halogen) is shorter than the sum of their Van der Waals radius. Depending on the structural and chemical environment, halogen bonding interactions can be weak or strong. In the case of some protein-ligand complexes, halogen bonds are energetically and geometrically comparable to that of hydrogen bonding if the donor-acceptor directionality remains consistent. This intermolecular interaction has been shown to be stabilizing and a conformational determinant in protein-ligand and DNA structures.
Chemical bonding of H2O
Water () is a simple triatomic bent molecule with C molecular symmetry and bond angle of 104.5° between the central oxygen atom and the hydrogen atoms. Despite being one of the simplest triatomic molecules, its chemical bonding scheme is nonetheless complex as many of its bonding properties such as bond angle, ionization energy, and electronic state energy cannot be explained by one unified bonding model. Instead, several traditional and advanced bonding models such as simple Lewis and VESPR structure, valence bond theory, molecular orbital theory, isovalent hybridization, and Bent's rule are discussed below to provide a comprehensive bonding model for , explaining and rationalizing the various electronic and physical properties and features manifested by its peculiar bonding arrangements.
Iodine
The melting and boiling points of iodine are the highest among the halogens, conforming to the increasing trend down the group, since iodine has the largest electron cloud among them that is the most easily polarised, resulting in its molecules having the strongest van der Waals interactions among the halogens. Similarly, iodine is the least volatile of the halogens. Because it has the largest atomic radius among the halogens, iodine has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the halogens.
The interhalogen bond in diiodine is the weakest of all the halogens. As such, 1% of a sample of gaseous iodine at atmospheric pressure is dissociated into iodine atoms at 575 °C. Temperatures greater than 750 °C are required for fluorine, chlorine, and bromine to dissociate to a similar extent. Most bonds to iodine are weaker than the analogous bonds to the lighter halogens. Gaseous iodine is composed of I molecules with an I–I bond length of 266.6 pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5 pm) in solid orthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbour xenon: the Xe–Xe bond length is 308.71 pm.) As such, within the iodine molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk iodine, to a shiny appearance and semiconducting properties. Iodine is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.
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
The dynamics of hydrogen bond structures in water can be probed by the IR spectrum of OH stretching vibration. 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. 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.
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.
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.
Hydrogen bond
The role of hydrogen bonds in protein folding has also been linked to osmolyte-induced protein stabilization. Protective osmolytes, such as trehalose and sorbitol, shift the protein folding equilibrium toward the folded state, in a concentration dependent manner. While the prevalent explanation for osmolyte action relies on excluded volume effects that are entropic in nature, recent circular dichroism (CD) experiments have shown osmolyte to act through an enthalpic effect. The molecular mechanism for their role in protein stabilization is still not well established, though several mechanisms have been proposed. Recently, computer molecular dynamics simulations suggested that osmolytes stabilize proteins by modifying the hydrogen bonds in the protein hydration layer.
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.
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.
GC-content
Qualitativelty, G (guanine) and C (cytosine) undergo a specific hydrogen bonding, whereas A (adenine) bonds specifically with T (thymine, in DNA) or U (uracil, in RNA). Quantitativelty, the GC pair is bound by three hydrogen bonds, while AT and AU pairs are bound by two hydrogen bonds. To emphasize this difference in the number of hydrogen bonds, the base pairings can be represented as respectively G≡C versus A=T and A=U. DNA with low GC-content is less stable than DNA with high GC-content; however, the hydrogen bonds themselves do not have a particularly significant impact on stabilization, the stabilization is due mainly to interactions of base stacking.
In spite of the higher thermostability conferred to the genetic material, it has been observed that at least some bacteria species with DNA of high GC-content undergo autolysis more readily, thereby reducing the longevity of the cell "per se". Due to the thermostability given to the genetic materials in high GC organisms, it was commonly believed that the GC content played a necessary role in adaptation temperatures, a hypothesis that was refuted in 2001.
However, it has been shown that there is a strong correlation between the prokaryotic optimal growth at higher temperatures and the GC content of structured RNAs (such as ribosomal RNA, transfer RNA, and many other non-coding RNAs). The AU base pairs are less stable than the GC base pairs previously attributed to GC bonds containing 3 hydrogen bonds and AU having only 2 hydrogen bonds, making high-GC-content RNA structures more resistant to the effects of high temperatures.
More recently, it has been proved that the most stabilizing factor of thermal stability of double stranded nucleic acids is actually due to the base stackings of adjacent bases, rather than the number of hydrogen bonds between the bases.There is more favorable stacking energy for G:C pairs because of the relative positions of exocyclic groups than in the A:U pairs. Additionally, there is a correlation between the order in which the bases stack and thermal stability.
Halogen bond
Halogen bonds are strong, specific, and directional interactions that give rise to well-defined structures. Halogen bond strengths range from 5–180 kJ/mol. The strength of XB allows it to compete with HB, which are a little bit weaker in strength. Halogen bonds tend to form at 180° angles, which was shown in Odd Hassel’s studies with bromine and 1,4-dioxane in 1954. Another contributing factor to halogen bond strength comes from the short distance between the halogen (Lewis acid, XB donor) and Lewis base (XB acceptor). The attractive nature of halogen bonds result in the distance between the donor and acceptor to be shorter than the sum of van der Waals radii. The XB interaction becomes stronger as the distance decreases between the halogen and Lewis base.
Kaonic hydrogen
Unlike in the hydrogen atom, where the binding between electron and proton is dominated by the electromagnetic interaction, kaons and protons interact also to a large extent by the strong interaction.
In kaonic hydrogen this strong contribution was found to be repulsive, shifting the ground state energy by 283 ± 36 (statistical) ± 6 (systematic) eV, thus making the system unstable with a resonance width of 541 ± 89 (stat) ± 22 (syst) eV (decay into Λπ and Σπ).
Bond-dissociation energy
The bond-dissociation energy (BDE, D", or DH°") is one measure of the strength of a chemical bond A–B. It can be defined as the standard enthalpy change when A–B is cleaved by homolysis to give fragments A and B, which are usually radical species. The enthalpy change is temperature dependent, and the bond-dissociation energy is often defined to be the enthalpy change of the homolysis at 0 K (absolute zero), although the enthalpy change at 298 K (standard conditions) is also a frequently encountered parameter. As a typical example, the bond-dissociation energy for one of the C–H bonds in ethane (CH) is defined as the standard enthalpy change of the process
Nonmetal
Hydrogen is a colourless, odourless, and comparatively unreactive diatomic gas with a density of 0.08988 × 10 g/cm and is about 14 times lighter than air. It condenses to a colourless liquid −252.879 °C and freezes into an ice- or snow-like solid at −259.16 °C. The solid form has a hexagonal crystalline structure and is soft and easily crushed. Hydrogen is an insulator in all of its forms. It has a high ionisation energy (1312.0 kJ/mol), moderate electron affinity (73 kJ/mol), and moderate electronegativity (2.2). Hydrogen is a poor oxidising agent (H + 2"e" → 2H = –2.25 V at pH 0). Its chemistry, most of which is based around its tendency to acquire the electron configuration of the noble gas helium, is largely covalent in nature, noting it can form ionic hydrides with highly electropositive metals, and alloy-like hydrides with some transition metals. The common oxide of hydrogen (HO) is a neutral oxide.
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.
Chemical bond
A single bond between two atoms corresponds to the sharing of one pair of electrons. The electron density of these two bonding electrons is concentrated in the region between the two atoms, which is the defining quality of a sigma bond.
A double bond between two atoms is formed by the sharing of two pairs of electrons, one in a sigma bond and one in a pi bond, with electron density concentrated on two opposite sides of the internuclear axis. A triple bond consists of three shared electron pairs, forming one sigma and two pi bonds.
Multipole density formalism
However, hydrogen atoms possess a feature unique among all the elements - they possess exactly one electron, which additionally is located on their valence shell and therefore is involved in creating strong covalent bonds with atoms of various other elements. While a bond is forming, the maximum of the electron density function moves significantly away from the nucleus and towards the other atom. This prevents any spherical approach from determining hydrogen position correctly by itself. Therefore, usually the hydrogen position is estimated basing on neutron crystallography data for similar molecules, or it is not modelled at all in the case of low-quality diffraction data.
Hydrogen storage
Most metal hydrides bind with hydrogen very strongly. As a result, high temperatures around 120 °C (248 °F) – 200 °C (392 °F) are required to release their hydrogen content. This energy cost can be reduced by using alloys which consists of a strong hydride former and a weak one such as in LiNH, LiBH and NaBH. These are able to form weaker bonds, thereby requiring less input to release stored hydrogen. However, if the interaction is too weak, the pressure needed for rehydriding is high, thereby eliminating any energy savings. The target for onboard hydrogen fuel systems is roughly <100 °C for release and <700 bar for recharge (20–60 kJ/mol H).
Dehalogenation
Among halogens, fluorine is the most electronegative atom and will have the highest tendency to make the strongest bond with carbon. The rate of dehalogenation depends on the bond strength between carbon and halogen atom. The bond dissociation energies of carbon-halogen bonds are described as: H3C-I (234 KJmol-1), H3C-Br (293 KJmol-1), H3C-Cl (351 KJmol-1), and H3C-F (452 KJmol-1). Thus, for the same structures the bond dissociation rate for dehalogenation will be: F« Cl Additionally, the rate of dehalogenation for alkyl halide also varies with steric environment and follows this trend: primary > secondary > tertiary halides.
Hydrogen clathrate
A hydrogen clathrate is a clathrate containing hydrogen in a water lattice. This substance is interesting due to its possible use to store hydrogen in a hydrogen economy. Another unusual characteristic is that multiple hydrogen molecules can occur at each cage site in the ice, one of only a very few guest molecule that forms clathrates with this property. The maximum ratio of hydrogen to water is 6 H to 17 HO. It can be formed at 250K in a diamond anvil at a pressure of 300MPa (3000 Bars). It takes about 30 minutes to form, so this method is impractical for rapid manufacture. The percent of weight of hydrogen is 3.77%. The cage compartments are hexakaidecahedral and hold from two to four molecules of hydrogen. At temperatures above 160K the molecules rotate around inside the cage. Below 120K the molecules stop racing around the cage, and below 50K are locked into a fixed position. This was determined with deuterium in a neutron scattering experiment.
Compounds of fluorine
Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant (pK) equal to 3.19. HF's weakness as an aqueous acid is paradoxical considering how polar the HF bond is, much more so than the bond in HCl, HBr, or HI. The explanation for the behavior is complicated, having to do with various cluster-forming tendencies of HF, water, and fluoride ion, as well as thermodynamic issues. At great concentrations, a property called homoconjugation is revealed. HF begins to accept fluoride ions, forming the polyatomic ions (such as bifluoride, ) and protons, thus greatly increasing the acidity of the compound. Hydrofluoric acid is also the strongest of the hydrohalic acids in acetic acid and similar solvents. Its hidden acidity potential is also revealed by the fact it protonates acids like hydrochloric, sulfuric, or nitric. Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass (hydrated only).
Solid nitrogen
Solid nitrogen has several properties relevant to its formation of rocks in the outer Solar System. 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. At 58 K the ultimate compressive strength is 0.24 MPa. Strength increases as temperature lowers becoming 0.54 MPa at 40.6 K. Elastic modulus varies from 161 to 225 MPa over the same range. Compared to other materials, solid nitrogen loses cohesion at low pressures and flows in the form of glaciers when amassed. Yet its density is higher than that of water ice, so the forces of buoyancy will naturally transport blocks of water ice towards the surface. This effect has been most clearly observed on Pluto (by the "New Horizons" space probe in 2015), where water ice makes up a major part of the surface layers as icebergs on top of nitrogen ice.
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.
Hydrogen bond
From interpretations of the anisotropies in the Compton profile of ordinary ice that the hydrogen bond is partly covalent. However, this interpretation was challenged.
Motorcycle boot
- Sole bond strength
Bond duration
where
Acid strength
The strength of a weak organic acid may depend on substituent effects. The strength of an inorganic acid is dependent on the oxidation state for the atom to which the proton may be attached. Acid strength is solvent-dependent. For example, hydrogen chloride is a strong acid in aqueous solution, but is a weak acid when dissolved in glacial acetic acid.
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.)
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).
Cohesion (chemistry)
Water, for example, is strongly cohesive as each molecule may make four hydrogen bonds to other water molecules in a tetrahedral configuration. This results in a relatively strong Coulomb force between molecules. In simple terms, the polarity (state in which a molecule is oppositely charged on its poles) of water molecules allows them to be attracted towards each other. The polarity is due to the electronegativity of the atom of oxygen: oxygen is more electronegative than the atoms of hydrogen, so the electrons they share through the covalent bonds are more often close to oxygen rather than hydrogen. These are called polar covalent bonds, covalent bonds between atoms that thus become oppositely charged. In the case of a water molecule, the hydrogen atoms carry positive charges while the oxygen atom has a negative charge. This charge polarization within the molecule allows it to align with adjacent molecules through strong intermolecular hydrogen bonding, rendering the bulk liquid cohesive. Van der Waals gases such as methane, however, have weak cohesion due only to van der Waals forces that operate by induced polarity in non-polar molecules.
Double bond
Double bonded compounds, alkene homologs, RE=ER are now known for all of the heavier group 14 elements. Unlike the alkenes these compounds are not planar but adopt twisted and/or trans bent structures. These effects become more pronounced for the heavier elements. The distannene (MeSi)CHSn=SnCH(SiMe) has a tin-tin bond length just a little shorter than a single bond, a trans bent structure with pyramidal coordination at each tin atom, and readily dissociates in solution to form (MeSi)CHSn: (stannanediyl, a carbene analog). The bonding comprises two weak donor acceptor bonds, the lone pair on each tin atom overlapping with the empty p orbital on the other. In contrast, in disilenes each silicon atom has planar coordination but the substituents are twisted so that the molecule as a whole is not planar. In diplumbenes the Pb=Pb bond length can be longer than that of many corresponding single bonds Plumbenes and stannenes generally dissociate in solution into monomers with bond enthalpies that are just a fraction of the corresponding single bonds. Some double bonds plumbenes and stannenes are similar in strength to hydrogen bonds. The Carter-Goddard-Malrieu-Trinquier model can be used to predict the nature of the bonding.
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.
Hydrogen bond
Many polymers are strengthened by hydrogen bonds within and between the chains. Among the synthetic polymers, a well characterized example is nylon, where hydrogen bonds occur in the repeat unit and play a major role in crystallization of the material. The bonds occur between carbonyl and amine groups in the amide repeat unit. They effectively link adjacent chains, which help reinforce the material. The effect is great in aramid fibre, where hydrogen bonds stabilize the linear chains laterally. The chain axes are aligned along the fibre axis, making the fibres extremely stiff and strong.
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.
Silicon monohydride
The Si-H bond strength is 80 kcal/mol.
Acid strength
While the p"K" value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water or DMSO), the tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weak aniline base) is measured by its Hammett acidity function, the "H" value. Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the p"K" and "H" values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water (p"K" = 3.2) or DMSO (p"K" = 15), has p"K" values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an "H" value of –15, making it a more strongly protonating medium than 100% sulfuric acid and thus, by definition, a superacid. (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its p"K" value (p"K" < –1.74). This usage is consistent with the common parlance of most practicing chemists.)
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
Hydrogen bonds are important in the structure of cellulose and derived polymers in its many different forms in nature, such as cotton and flax.
Chalcogen bond
A chalcogen bond is an attractive interaction in the family of σ-hole interactions, along with hydrogen bonds and halogen bonds. This family of attractive interactions has been modeled as an electron donor interacting with the σ* orbital of a C-X bond (X= hydrogen, halogen, chalcogen, pnictogen, etc.). Electron density mapping is often invoked to visualize the electron density of the donor and an electrophilic region on the acceptor, referred to as a σ-hole. Chalcogen bonds, much like hydrogen and halogen bonds, have been invoked in various non-covalent interactions, such as protein folding, crystal engineering, self-assembly, catalysis, transport, sensing, templation, and drug design.
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.
Bond valence method
Since the bonding strength of an atom is the valence expected for a bond formed by that atom, it follows that the most stable bonds will be formed between atoms with the same bonding strengths. In practice some tolerance is allowed, but bonds are rarely formed if the ratio of the bonding strengths of the two atoms exceeds two, a condition expressed by the inequality shown in Eq. 4. This is known and the "valence matching rule".
Carbon–carbon bond
Relative to most bonds, a carbon–carbon bond is very strong.
Bond energy
In chemistry, bond energy ("E") or bond enthalpy ("H") is the measure of bond strength in a chemical bond. IUPAC defines bond energy as the average value of the gas-phase bond dissociation energies (usually at a temperature of 298 K) for all bonds of the same type within the same chemical species. For example, the carbon–hydrogen bond energy in methane "H"(C–H) is the enthalpy change involved with breaking up one molecule of methane into a carbon atom and four hydrogen radicals, divided by 4. Tabulated bond energies are generally values of bond energies averaged over a number of selected typical chemical species containing that type of bond. Bond energy ("E") or bond enthalpy ("H") should not be confused with bond-dissociation energy. Bond energy is the average of all the bond-dissociation energies in a molecule, and will show a different value for a given bond than the bond-dissociation energy would. This is because the energy required to break a single bond in a specific molecule differs for each bond in that molecule. For example, methane has four C–H bonds and the bond-dissociation energies are 435 kJ/mol for "D"(CH–H), 444 kJ/mol for "D"(CH–H), 444 kJ/mol for "D"(CH–H) and 339 kJ/mol for "D"(C–H). Their average, and hence the bond energy, is 414 kJ/mol, even though not a single bond required specifically 414 kJ/mol to be broken.
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.
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
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.
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
Halogen bond
A parallel relationship can easily be drawn between halogen bonding and hydrogen bonding (HB). In both types of bonding, an electron donor/electron acceptor relationship exists. The difference between the two is what species can act as the electron donor/electron acceptor. In hydrogen bonding, a hydrogen atom acts as the electron acceptor and forms a non-covalent interaction by accepting electron density from an electron rich site (electron donor). In halogen bonding, a halogen atom is the electron acceptor. Simultaneously, the normal covalent bond between H or X and D weakens, so the electron density on H or X appears to be reduced. Electron density transfers results in a penetration of the van der Waals volumes.
Bond hardening
The second pair of proton peaks (1 eV KER) comes from bond softening of the H ion, which dissociates into a proton and a neutral hydrogen atom (undetected). The dissociation starts at the 3-photon gap and proceeds to the 2ω limit (the lower blue arrow in Fig. 1). Since both the initial and the final energies of this process are fixed by the 1.55 eV photon energy, the KER is also constant producing the two vertical lines in Fig. 2.
Carbon–hydrogen bond
The carbon-hydrogen bond (C–H bond) is a bond between carbon and hydrogen atoms that can be found in many organic compounds. This bond is a covalent bond meaning that carbon shares its outer valence electrons with up to four hydrogens. This completes both of their outer shells making them stable. Carbon–hydrogen bonds have a bond length of about 1.09 Å (1.09 × 10 m) and a bond energy of about 413 kJ/mol (see table below). Using Pauling's scale—C (2.55) and H (2.2)—the electronegativity difference between these two atoms is 0.35. Because of this small difference in electronegativities, the C−H bond is generally regarded as being non-polar. In structural formulas of molecules, the hydrogen atoms are often omitted. Compound classes consisting solely of C–H bonds and C–C bonds are alkanes, alkenes, alkynes, and aromatic hydrocarbons. Collectively they are known as hydrocarbons.
PH
Strong acids and bases are compounds that, for practical purposes, are completely dissociated in water. Under normal circumstances this means that the concentration of hydrogen ions in acidic solution can be taken to be equal to the concentration of the acid. The pH is then equal to minus the logarithm of the concentration value. Hydrochloric acid (HCl) is an example of a strong acid. The pH of a 0.01M solution of HCl is equal to −log(0.01), that is, pH = 2. Sodium hydroxide, NaOH, is an example of a strong base. The p[OH] value of a 0.01M solution of NaOH is equal to −log(0.01), that is, p[OH] = 2. From the definition of p[OH] above, this means that the pH is equal to about 12. For solutions of sodium hydroxide at higher concentrations the self-ionization equilibrium must be taken into account.
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).
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.
Aurophilicity
The main evidence for aurophilicity is from the crystallographic analysis of Au(I) complexes. The aurophilic bond has a length of about 3.0 Å and a strength of about 7–12 kcal/mol, which is comparable to the strength of a hydrogen bond. The effect is greatest for gold as compared to copper or silver—the higher elements in its periodic table group—due to increased relativistic effects. Observations and theory show that, on average, 28% of the binding energy in the aurophilic interaction can be attributed to relativistic expansion of the gold d orbitals.
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.
Dihydrogen bond
In chemistry, a dihydrogen bond is a kind of hydrogen bond, an interaction between a metal hydride bond and an OH or NH group or other proton donor. With a van der Waals radius of 1.2 Å, hydrogen atoms do not usually approach other hydrogen atoms closer than 2.4 Å. Close approaches near 1.8 Å, are, however, characteristic of dihydrogen bonding.
An early example of this phenomenon is credited to Brown and Heseltine. They observed intense absorptions in the IR bands at 3300 and 3210 cm for a solution of (CH)NHBH. The higher energy band is assigned to a normal N−H vibration whereas the lower energy band is assigned to the same bond, which is interacting with the B−H. Upon dilution of the solution, the 3300 cm band increased in intensity and the 3210 cm band decreased, indicative of intermolecular association.
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.
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.
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.
Metal ions in aqueous solution
The strength of the metal-oxygen bond can be estimated in various ways. The hydration enthalpy, though based indirectly on experimental measurements, is the most reliable measure. The scale of values is based on an arbitrarily chosen zero, but this does not affect differences between the values for two metals. Other measures include the M–O vibration frequency and the M–O bond length. The strength of the M-O bond tends to increase with the charge and decrease as the size of the metal ion increases. In fact there is a very good linear correlation between hydration enthalpy and the ratio of charge squared to ionic radius, z/r. For ions in solution Shannon's "effective ionic radius" is the measure most often used.
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.
Hydrogen bond
The hydrogen bond can be compared with the closely related dihydrogen bond, which is also an intermolecular bonding interaction involving hydrogen atoms. These structures have been known for some time, and well characterized by crystallography; however, an understanding of their relationship to the conventional hydrogen bond, ionic bond, and covalent bond remains unclear. Generally, the hydrogen bond is characterized by a proton acceptor that is a lone pair of electrons in nonmetallic atoms (most notably in the nitrogen, and chalcogen groups). In some cases, these proton acceptors may be pi-bonds or metal complexes. In the dihydrogen bond, however, a metal hydride serves as a proton acceptor, thus forming a hydrogen-hydrogen interaction. Neutron diffraction has shown that the molecular geometry of these complexes is similar to hydrogen bonds, in that the bond length is very adaptable to the metal complex/hydrogen donor system.
Iron–hydrogen alloy
The density deficit is thought to be due to mixture of lighter elements such as silicon or carbon. Hydrogen has been thought unlikely because of its volatility, but recent studies have uncovered plausible mechanisms for its incorporation and permanence in the core. It is estimated that hcp FeH would be stable under those conditions. Iron–hydrogen alloys could have been formed in a reaction of iron with water in magma during the formation of the earth. Above 5 GPa, iron will split water yielding the hydride and ferrous ions:
Indeed, Okuchi obtained magnetite and iron hydride by reacting magnesium silicate, magnesium oxide, silica and water with metallic iron in a diamond cell at 2000 C. Okuchi argues that most of the hydrogen accreted to Earth should have dissolved into the primeval magma ocean; and if the pressure at the bottom of the magma was 7.5 GPa or more, then almost all of that hydrogen would have reacted with iron to form the hydride, which then would have sunk to the core where it would be stabilized by the increased pressure. Moreover, it appears that at those pressures iron binds hydrogen in preference to carbon.
Hydrogen bond
The number of hydrogen bonds formed by a molecule of liquid water fluctuates with time and temperature. From TIP4P liquid water simulations at 25 °C, it was estimated that each water molecule participates in an average of 3.59 hydrogen bonds. At 100 °C, this number decreases to 3.24 due to the increased molecular motion and decreased density, while at 0 °C, the average number of hydrogen bonds increases to 3.69. A more recent study found a much smaller number of hydrogen bonds: 2.357 at 25 °C. The differences may be due to the use of a different method for defining and counting the hydrogen bonds.
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
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.
Sigma bond
In chemistry, sigma bonds (σ bonds) are the strongest type of covalent chemical bond. They are formed by head-on overlapping between atomic orbitals. Sigma bonding is most simply defined for diatomic molecules using the language and tools of symmetry groups. In this formal approach, a σ-bond is symmetrical with respect to rotation about the bond axis. By this definition, common forms of sigma bonds are s+s, p+p, s+p and d+d (where z is defined as the axis of the bond or the internuclear axis).
Quantum theory also indicates that molecular orbitals (MO) of identical symmetry actually mix or "hybridize". As a practical consequence of this mixing of diatomic molecules, the wavefunctions s+s and p+p molecular orbitals become blended. The extent of this mixing (or hybridization or blending) depends on the relative energies of the MOs of like symmetry.
For homodiatomics, bonding σ orbitals have no nodal planes at which the wavefunction is zero, either between the bonded atoms or passing through the bonded atoms. The corresponding antibonding, or σ* orbital, is defined by the presence of one nodal plane between the two bonded atoms.
Oxyacid
All oxyacids have the acidic hydrogen bound to an oxygen atom, so bond strength (length) is not a factor, as it is with binary nonmetal hydrides. Rather, the electronegativity of the central atom (X) and the number of O atoms determine oxyacid acidity. With the same central atom X, acid strength increases as the number of oxygens attached to X increases. With the same number of oxygens around E, acid strength increases with the electronegativity of X.
Hydrogen bond
Wool, being a protein fibre, is held together by hydrogen bonds, causing wool to recoil when stretched. However, washing at high temperatures can permanently break the hydrogen bonds and a garment may permanently lose its shape.
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.
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.
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).
Carbon–fluorine bond
The high electronegativity of fluorine (4.0 for F vs. 2.5 for carbon) gives the carbon–fluorine bond a significant polarity/dipole moment. The electron density is concentrated around the fluorine, leaving the carbon relatively electron poor. This introduces ionic character to the bond through partial charges (C—F). The partial charges on the fluorine and carbon are attractive, contributing to the unusual bond strength of the carbon–fluorine bond. The bond is labeled as "the strongest in organic chemistry," because fluorine forms the strongest single bond to carbon. Carbon–fluorine bonds can have a bond dissociation energy (BDE) of up to 544 kJ/mol. The BDE (strength of the bond) is higher than other carbon–halogen and carbon–hydrogen bonds. For example, the molecule represented by CHX has a BDE of 115 kcal/mol for carbon–fluorine while values of 104.9, 83.7, 72.1, and 57.6 kcal/mol represent carbon–X bonds to hydrogen, chlorine, bromine, and iodine, respectively.
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.
Pi bond
Pi bonds are usually weaker than sigma bonds. The C-C double bond, composed of one sigma and one pi bond, has a bond energy less than twice that of a C-C single bond, indicating that the stability added by the pi bond is less than the stability of a sigma bond. From the perspective of quantum mechanics, this bond's weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation. This is contrasted by sigma bonds which form bonding orbitals directly between the nuclei of the bonding atoms, resulting in greater overlap and a strong sigma bond.
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.
Hydride
Aside from electride, the hydride ion is the simplest possible anion, consisting of two electrons and a proton. Hydrogen has a relatively low electron affinity, 72.77 kJ/mol and reacts exothermically with protons as a powerful Lewis base. 
The low electron affinity of hydrogen and the strength of the H–H bond (∆H = 436 kJ/mol) means that the hydride ion would also be a strong reducing agent
Hydronium
A 2007 calculation of the enthalpies and free energies of the various hydrogen bonds around the hydronium cation in liquid protonated water at room temperature and a study of the proton hopping mechanism using molecular dynamics showed that the hydrogen-bonds around the hydronium ion (formed with the three water ligands in the first solvation shell of the hydronium) are quite strong compared to those of bulk water.
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.
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.