Instruct: Given a question, retrieve documents that can help answer the question.
Query: What is the strength of a hydrogen bond?
Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule).[4] 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.[5] 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.
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.
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.
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;[38] 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.[38]
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.
The intermolecular forces in polymers can be affected by dipoles in the monomer units. Polymers containing amide or carbonyl groups can form hydrogen bonds between adjacent chains; the partially positively charged hydrogen atoms in N-H groups of one chain are strongly attracted to the partially negatively charged oxygen atoms in C=O groups on another. These strong hydrogen bonds, for example, result in the high tensile strength and melting point of polymers containing urethane or urea linkages. Polyesters have dipole-dipole bonding between the oxygen atoms in C=O groups and the hydrogen atoms in H-C groups. Dipole bonding is not as strong as hydrogen bonding, so a polyester's melting point and strength are lower than Kevlar's (Twaron), but polyesters have greater flexibility.
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").
In contrast to this trend, it is understood that an increasingly negative Δ"V" is correlated with an increase of strength in hydrogen bonding, which is the opposing trend for the above; this indicates that the 2"β"-substituents (at least for the heterocyclic class) are dominated by electrostatic factors for binding in-the-stead of the presumptive hydrogen bonding model for this substituent of the cocaine-like binding ligand.
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 3-center 4-electron bond. This type of bond is much stronger than "normal" hydrogen bonds, in fact, its strength is comparable to a covalent bond. It is seen in ice at high pressure (Ice X), 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]. Much has been done to explain the symmetric hydrogen bond quantum-mechanically, as it seems to violate the duet rule for the first shell: The proton is effectively surrounded by four electrons. Because of this problem, some consider it to be an ionic bond.
The substrate, adenosine, is stabilized and bound to the active site by nine hydrogen bonds. The carboxyl group of Glu217, roughly coplanar with the substrate purine ring, is in position to form a hydrogen bond with N1 of the substrate. The carboxyl group of Asp296, also coplanar with the substrate purine ring, forms hydrogen bond with N7 of the substrate. The NH group of Gly184 is in position to form a hydrogen bond with N3 of the substrate. Asp296 forms bonds both with the Zn ion as well as with 6-OH of the substrate. His238 also hydrogen bonds to substrate 6-OH. The 3'-OH of the substrate ribose forms a hydrogen bond with Asp19, while the 5'-OH forms a hydrogen bond with His17. Two further hydrogen bonds are formed to water molecules, at the opening of the active site, by the 2'-OH and 3'-OH of the substrate.
There is debate about the effectiveness of methane and other hydrocarbons as a solvent for life compared to water or ammonia. Water is a stronger solvent than the hydrocarbons, enabling easier transport of substances in a cell. However, water is also more chemically reactive, and can break down large organic molecules through hydrolysis. A life-form whose solvent was a hydrocarbon would not face the threat of its biomolecules being destroyed in this way. Also, the water molecule's tendency to form strong hydrogen bonds can interfere with internal hydrogen bonding in complex organic molecules. Life with a hydrocarbon solvent could make more use of hydrogen bonds within its biomolecules. Moreover, the strength of hydrogen bonds within biomolecules would be appropriate to a low-temperature biochemistry.
Wrinkles form in cotton fabrics due to the free hydroxyl groups. Cotton is a form of cellulose chains linked to form firm three-dimensional structures that offer both tensile strength and flexibility due to their carbon-carbon and carbon-oxygen bond based backbone. Since cellulose is composed of glucose units, cyclic carbohydrate molecules, cellulose has free hydroxyl groups (-OH) projecting from each monomeric subunit. These hydroxyl groups tend form hydrogen-bonds to neighboring hydroxyl groups. When the fabric is stressed either by heat or pressure, the original hydrogen bonds in the cotton fabric break and reform with nearby atoms at random. This re-forming of the hydrogen bonds, known as cross linking, is the reason for wrinkles, or creases, in the fabric. DMEU works to prevent these wrinkles by covalently bonding to two free hydroxyl groups in the fabric through a dehydration reaction that is not as easily broken as the hydrogen bond before treatment.
As seen in the figure, this formation of a C-O bond through a dehydration reaction (loss of water) allows the fabric to be bleached and heated to reasonable temperatures without fear of breaking the C-O bonds formed.
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 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.
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.
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]−.
Factors affecting the rate of olefin insertions include the formation of the cyclic, planar, four-center transition state involving incipient formation of a bond between the metal and an olefin carbon. From this transition state, it can be seen that a partial positive charge forms on the β-carbon with a partial negative charge formed on the carbon initially bonded to the metal. This polarization explains the subsequently observed formation of the bond between the negatively charged carbon/hydrogen and the positively charged β-carbon as well as the simultaneous formation of the metal-α-carbon bond. This transition state also highlights the two factors that most strongly contribute to the rate of olefin insertion reactions: (i) orbital overlap of the alkyl group initially attached to the metal and (ii) the strength of the metal-alkyl bond. With greater orbital overlap between the partially positive β-carbon and the partially negative hydrogen/alkyl group carbon, the formation of the new C-C bond is facilitated. With increasing strength of the metal-alkyl bond, the breaking of the bond between the metal and the hydrogen/alkyl carbon bond to form the two new bonds with the α-carbon and β-carbon (respectively) is slower, thus decreasing the rate of the insertion reaction.
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.
forming a total of three sp orbitals with one remaining p orbital. In ethylene (ethene) the two carbon atoms form a σ bond by overlapping one sp orbital from each carbon atom. The π bond between the carbon atoms perpendicular to the molecular plane is formed by 2p–2p overlap. Each carbon atom forms covalent C–H bonds with two hydrogens by s–sp overlap, all with 120° bond angles. The hydrogen–carbon bonds are all of equal strength and length, in agreement with experimental data.
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.
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.
A hydrogen atom attached to a relatively electronegative atom is the hydrogen bond donor.[9] C-H bonds only participate in hydrogen bonding when the carbon atom is bound to electronegative substituents, as is the case in chloroform, CHCl3.[10] In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named proton acceptor, whereas the one covalently bound to the hydrogen is named the proton donor. In the donor molecule, the H center is protic. The donor is a Lewis base. Hydrogen bonds are represented as H···Y system, where the dots represent the hydrogen bond. Liquids that display hydrogen bonding (such as water) are called associated liquids.
Because water may form hydrogen bonds with solute proton donors and acceptors, it may competitively inhibit the formation of solute intermolecular or intramolecular hydrogen bonds. Consequently, hydrogen bonds between or within solute molecules dissolved in water are almost always unfavorable relative to hydrogen bonds between water and the donors and acceptors for hydrogen bonds on those solutes.[28] Hydrogen bonds between water molecules have an average lifetime of 10−11 seconds, or 10 picoseconds.[29]
Hydrogen spillover increases with adsorption temperature and metal dispersion. A correlation has been reported between available surface area and the capacity for hydrogen storage. For PdnP-containing MOFs, in the presence of saturated metal particles, the capacity for hydrogen spillover only relied on the sorbent’s surface area and pore size. On catalysts such as platinum or nickel, atomic hydrogen can be generated at a high frequency. Through surface diffusion, multi-functional transport of hydrogen atoms can enhance a reaction and even regenerate a catalyst. However, problems present in the strength of the hydrogen-support bond; too strong of an interaction would hinder its extraction via reverse spillover and nullify its function as a fuel cell. Conversely, too weak a bond and the hydrogens are easily lost to the environment.
Hydrogen bonds can vary in strength from weak (1–2 kJ mol−1) to strong (161.5 kJ mol−1 in the ion HF− 2).[11][12] Typical enthalpies in vapor include: