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102-772: Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule). The energy of a hydrogen bond depends on the geometry, the environment, and the nature of the specific donor and acceptor atoms and 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. Hydrogen bonds are responsible for holding materials such as paper and felted wool together, and for causing separate sheets of paper to stick together after becoming wet and subsequently drying. The hydrogen bond

204-465: A {\displaystyle a} to point b {\displaystyle b} with the following line integral : From these equations, we see that the electric potential is constant in any region for which the electric field vanishes (such as occurs inside a conducting object). A test particle 's potential energy, U E single {\displaystyle U_{\mathrm {E} }^{\text{single}}} , can be calculated from

306-408: A line integral of the work, q n E ⋅ d ℓ {\displaystyle q_{n}\mathbf {E} \cdot \mathrm {d} \mathbf {\ell } } . We integrate from a point at infinity, and assume a collection of N {\displaystyle N} particles of charge Q n {\displaystyle Q_{n}} , are already situated at

408-583: A triple integral : Gauss's law states that "the total electric flux through any closed surface in free space of any shape drawn in an electric field is proportional to the total electric charge enclosed by the surface." Many numerical problems can be solved by considering a Gaussian surface around a body. Mathematically, Gauss's law takes the form of an integral equation: where d 3 r = d x   d y   d z {\displaystyle \mathrm {d} ^{3}r=\mathrm {d} x\ \mathrm {d} y\ \mathrm {d} z}

510-474: A Keesom interaction depends on the inverse sixth power of the distance, unlike the interaction energy of two spatially fixed dipoles, which depends on the inverse third power of the distance. The Keesom interaction can only occur among molecules that possess permanent dipole moments, i.e., two polar molecules. Also Keesom interactions are very weak van der Waals interactions and do not occur in aqueous solutions that contain electrolytes. The angle averaged interaction

612-639: A big role with this. Concerning electron density topology, recent methods based on electron density gradient methods have emerged recently, notably with the development of IBSI (Intrinsic Bond Strength Index), relying on the IGM (Independent Gradient Model) methodology. Electrostatic Electrostatics is a branch of physics that studies slow-moving or stationary electric charges . Since classical times , it has been known that some materials, such as amber , attract lightweight particles after rubbing . The Greek word for amber, ἤλεκτρον ( ḗlektron ),

714-692: A donor, particularly when the carbon or one of its neighbors is electronegative (e.g., in chloroform, aldehydes and terminal acetylenes). Gradually, it was recognized that there are many examples of weaker hydrogen bonding involving donor other than N, O, or F and/or acceptor Ac with electronegativity approaching that of hydrogen (rather than being much more electronegative). Although weak (≈1 kcal/mol), "non-traditional" hydrogen bonding interactions are ubiquitous and influence structures of many kinds of materials. The definition of hydrogen bonding has gradually broadened over time to include these weaker attractive interactions. In 2011, an IUPAC Task Group recommended

816-582: A feat that would only be possible if the hydrogen bond contained some covalent character. The concept of hydrogen bonding once was challenging. Linus Pauling credits T. S. Moore and T. F. Winmill with the first mention of the hydrogen bond, in 1912. Moore and Winmill used the hydrogen bond to account for the fact that trimethylammonium hydroxide is a weaker base than tetramethylammonium hydroxide . The description of hydrogen bonding in its better-known setting, water, came some years later, in 1920, from Latimer and Rodebush. In that paper, Latimer and Rodebush cited

918-407: A fundamental, unifying theory that is able to explain the various types of interactions such as hydrogen bonding , van der Waals force and dipole–dipole interactions. Typically, this is done by applying the ideas of quantum mechanics to molecules, and Rayleigh–Schrödinger perturbation theory has been especially effective in this regard. When applied to existing quantum chemistry methods, such

1020-501: A hydrogen each, forming two additional hydrogen bonds, and the second hydrogen atom also interacts with a neighbouring oxygen. Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides , which have little capability to hydrogen bond. Intramolecular hydrogen bonding is partly responsible for the secondary , tertiary , and quaternary structures of proteins and nucleic acids . It also plays an important role in

1122-522: A large number of electrons will have a greater associated London force than an atom with fewer electrons. The dispersion (London) force is the most important component because all materials are polarizable, whereas Keesom and Debye forces require permanent dipoles. The London interaction is universal and is present in atom-atom interactions as well. For various reasons, London interactions (dispersion) have been considered relevant for interactions between macroscopic bodies in condensed systems. Hamaker developed

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1224-451: A modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry . This definition specifies: The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X−H in which X is more electronegative than H, and an atom or a group of atoms in the same or another molecule, in which there

1326-475: A net attraction between them. Examples of polar molecules include hydrogen chloride (HCl) and chloroform (CHCl 3 ). Often molecules contain dipolar groups of atoms, but have no overall dipole moment on the molecule as a whole. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane and carbon dioxide . The dipole–dipole interaction between two individual atoms

1428-414: A quantum mechanical explanation of intermolecular interactions provides an array of approximate methods that can be used to analyze intermolecular interactions. One of the most helpful methods to visualize this kind of intermolecular interactions, that we can find in quantum chemistry, is the non-covalent interaction index , which is based on the electron density of the system. London dispersion forces play

1530-404: A solid or liquid, i.e., a condensed phase. Lower temperature favors the formation of a condensed phase. In a condensed phase, there is very nearly a balance between the attractive and repulsive forces. Intermolecular forces observed between atoms and molecules can be described phenomenologically as occurring between permanent and instantaneous dipoles, as outlined above. Alternatively, one may seek

1632-447: A substrate and an enzyme or a molecule with a catalyst , but several such weak interactions with the required spatial configuration of the active center of the enzyme lead to significant restructuring changes the energy state of molecules or substrate, which ultimately leads to the breaking of some and the formation of other covalent chemical bonds. Strictly speaking, all enzymatic reactions begin with intermolecular interactions between

1734-412: A water molecule is up to four. 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,

1836-452: A way that resembles integration by parts . These two integrals for electric field energy seem to indicate two mutually exclusive formulas for electrostatic energy density, namely 1 2 ρ ϕ {\textstyle {\frac {1}{2}}\rho \phi } and 1 2 ε 0 E 2 {\textstyle {\frac {1}{2}}\varepsilon _{0}E^{2}} ; they yield equal values for

1938-527: A weakening of the X−H bond. Certain hydrogen bonds - improper hydrogen bonds - show a blue shift of the X−H stretching frequency and a decrease in the bond length. H-bonds can also be measured by IR vibrational mode shifts of the acceptor. The amide I mode of backbone carbonyls in α-helices shifts to lower frequencies when they form H-bonds with side-chain hydroxyl groups. The dynamics of hydrogen bond structures in water can be probed by this OH stretching vibration. In

2040-454: 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. Liquid water's high boiling point

2142-412: Is a unit vector that indicates the direction of the field. For a single point charge, q {\displaystyle q} , at the origin, the magnitude of this electric field is E = q / 4 π ε 0 r 2 {\displaystyle E=q/4\pi \varepsilon _{0}r^{2}} and points away from that charge if it is positive. The fact that

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2244-420: Is a vector field that can be defined everywhere, except at the location of point charges (where it diverges to infinity). It is defined as the electrostatic force , {\displaystyle \mathbf {,} } on a hypothetical small test charge at the point due to Coulomb's law, divided by the charge q {\displaystyle q} Electric field lines are useful for visualizing

2346-427: 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

2448-432: Is a pair of water molecules with one hydrogen bond between them, which 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

2550-421: Is a strong type of hydrogen bond. It is characterized by the π-delocalization that involves the hydrogen and cannot be properly described by the electrostatic model alone. This description of the hydrogen bond has been proposed to describe unusually short distances generally observed between O=C−OH··· or ···O=C−C=C−OH . The X−H distance is typically ≈110  pm , whereas

2652-556: Is a volume element. If the charge is distributed over a surface or along a line, replace ρ d 3 r {\displaystyle \rho \,\mathrm {d} ^{3}r} by σ d A {\displaystyle \sigma \,\mathrm {d} A} or λ d ℓ {\displaystyle \lambda \,\mathrm {d} \ell } . The divergence theorem allows Gauss's Law to be written in differential form: where ∇ ⋅ {\displaystyle \nabla \cdot }

2754-406: Is also responsible for many of the physical and chemical properties of compounds of N, O, and F that seem unusual compared with other similar structures. In particular, 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

2856-414: Is also seen in the bifluoride ion [F···H···F] . Due to severe steric constraint, the protonated form of Proton Sponge (1,8-bis(dimethylamino)naphthalene) and its derivatives also have symmetric hydrogen bonds ( [N···H···N] ), although in the case of protonated Proton Sponge, the assembly is bent. The hydrogen bond can be compared with

2958-408: 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. 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

3060-587: Is attractive. If r {\displaystyle r} is the distance (in meters ) between two charges, then the force between two point charges Q {\displaystyle Q} and q {\displaystyle q} is: where ε 0 = 8.854 187 8188 (14) × 10  F⋅m ‍ is the vacuum permittivity . The SI unit of ε 0 is equivalently A ⋅ s ⋅kg ⋅m or C ⋅ N ⋅m or F ⋅m . The electric field, E {\displaystyle \mathbf {E} } , in units of Newtons per Coulomb or volts per meter,

3162-728: Is called the Debye force , named after Peter J. W. Debye . One example of an induction interaction between permanent dipole and induced dipole is the interaction between HCl and Ar. In this system, Ar experiences a dipole as its electrons are attracted (to the H side of HCl) or repelled (from the Cl side) by HCl. The angle averaged interaction is given by the following equation: where α 2 {\displaystyle \alpha _{2}} = polarizability. This kind of interaction can be expected between any polar molecule and non-polar/symmetrical molecule. The induction-interaction force

Hydrogen bond - Misplaced Pages Continue

3264-416: 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

3366-485: Is equal to the number of active pairs. The molecule which donates its hydrogen is termed the donor molecule, while the molecule containing lone pair participating in H bonding is termed the acceptor molecule. The number of active pairs is equal to the common number between number of hydrogens the donor has and the number of lone pairs the acceptor has. Though both not depicted in the diagram, water molecules have four active bonds. The oxygen atom’s two lone pairs interact with

3468-551: Is evidence of bond formation. Hydrogen bonds can vary in strength from weak (1–2 kJ/mol) to strong (161.5 kJ/mol in the bifluoride ion, HF − 2 ). Typical enthalpies in vapor include: The strength of intermolecular hydrogen bonds is most often evaluated by measurements of equilibria between molecules containing donor and/or acceptor units, most often in solution. The strength of intramolecular hydrogen bonds can be studied with equilibria between conformers with and without hydrogen bonds. The most important method for

3570-603: Is far weaker than dipole–dipole interaction, but stronger than the London dispersion force . The third and dominant contribution is the dispersion or London force (fluctuating dipole–induced dipole), which arises due to the non-zero instantaneous dipole moments of all atoms and molecules. Such polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in non-polar molecules. Thus, London interactions are caused by random fluctuations of electron density in an electron cloud. An atom with

3672-473: Is formed. When the spacing is less, between positions i and i + 3 , then a 3 10 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 ). Bifurcated H-bond systems are common in alpha-helical transmembrane proteins between

3774-489: Is given by the following equation: where d = electric dipole moment, ε 0 {\displaystyle \varepsilon _{0}} = permittivity of free space, ε r {\displaystyle \varepsilon _{r}} = dielectric constant of surrounding material, T = temperature, k B {\displaystyle k_{\text{B}}} = Boltzmann constant, and r = distance between molecules. The second contribution

3876-407: Is more important depends on temperature and pressure (see compressibility factor ). In a gas, the distances between molecules are generally large, so intermolecular forces have only a small effect. The attractive force is not overcome by the repulsive force, but by the thermal energy of the molecules. Temperature is the measure of thermal energy, so increasing temperature reduces the influence of

3978-652: Is much stronger than the forces present between neighboring molecules. Both sets of forces are essential parts of force fields frequently used in molecular mechanics . The first reference to the nature of microscopic forces is found in Alexis Clairaut 's work Théorie de la figure de la Terre, published in Paris in 1743. Other scientists who have contributed to the investigation of microscopic forces include: Laplace , Gauss , Maxwell , Boltzmann and Pauling . Attractive intermolecular forces are categorized into

4080-505: Is not so for big moving systems like enzyme molecules interacting with substrate molecules. Here the numerous intramolecular (most often - hydrogen bonds ) bonds form an active intermediate state where the intermolecular bonds cause some of the covalent bond to be broken, while the others are formed, in this way proceeding the thousands of enzymatic reactions , so important for living organisms . Intermolecular forces are repulsive at short distances and attractive at long distances (see

4182-447: 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 their van der Waals radii , and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence . The number of Hydrogen bonds formed between molecules

Hydrogen bond - Misplaced Pages Continue

4284-467: Is still not well established, though several mechanisms have been proposed. Computer molecular dynamics simulations suggest that osmolytes stabilize proteins by modifying the hydrogen bonds in the protein hydration layer. 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

4386-492: Is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. These interactions tend to align the molecules to increase attraction (reducing potential energy ). An example of a dipole–dipole interaction can be seen in hydrogen chloride (HCl): the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have

4488-412: Is the divergence operator . The definition of electrostatic potential, combined with the differential form of Gauss's law (above), provides a relationship between the potential Φ and the charge density ρ : This relationship is a form of Poisson's equation . In the absence of unpaired electric charge, the equation becomes Laplace's equation : The validity of the electrostatic approximation rests on

4590-612: Is the Lewis base. Hydrogen bonds are represented as H···Y system, where the dots represent the hydrogen bond. Liquids that display hydrogen bonding (such as water) are called associated liquids . Hydrogen bonds arise from a combination of electrostatics (multipole-multipole and multipole-induced multipole interactions), covalency (charge transfer by orbital overlap), and dispersion ( London forces ). In weaker hydrogen bonds, hydrogen atoms tend to bond to elements such as sulfur (S) or chlorine (Cl); even carbon (C) can serve as

4692-408: Is the force that mediates interaction between molecules, including the electromagnetic forces of attraction or repulsion which act between atoms and other types of neighbouring particles, e.g. atoms or ions . Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. For example, the covalent bond , involving sharing electron pairs between atoms,

4794-567: Is the induction (also termed polarization) or Debye force, arising from interactions between rotating permanent dipoles and from the polarizability of atoms and molecules (induced dipoles). These induced dipoles occur when one molecule with a permanent dipole repels another molecule's electrons. A molecule with permanent dipole can induce a dipole in a similar neighboring molecule and cause mutual attraction. Debye forces cannot occur between atoms. The forces between induced and permanent dipoles are not as temperature dependent as Keesom interactions because

4896-445: Is usually zero, since atoms rarely carry a permanent dipole. The Keesom interaction is a van der Waals force. It is discussed further in the section "Van der Waals forces". Ion–dipole and ion–induced dipole forces are similar to dipole–dipole and dipole–induced dipole interactions but involve ions, instead of only polar and non-polar molecules. Ion–dipole and ion–induced dipole forces are stronger than dipole–dipole interactions because

4998-418: Is what would be measured at r i {\displaystyle \mathbf {r} _{i}} if the charge Q i {\displaystyle Q_{i}} were missing. This formula obviously excludes the (infinite) energy that would be required to assemble each point charge from a disperse cloud of charge. The sum over charges can be converted into an integral over charge density using

5100-468: 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,

5202-473: The Lennard-Jones potential ). In a gas, the repulsive force chiefly has the effect of keeping two molecules from occupying the same volume. This gives a real gas a tendency to occupy a larger volume than an ideal gas at the same temperature and pressure. The attractive force draws molecules closer together and gives a real gas a tendency to occupy a smaller volume than an ideal gas. Which interaction

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5304-403: The electrostatic potential (also known as the voltage ). An electric field, E {\displaystyle E} , points from regions of high electric potential to regions of low electric potential, expressed mathematically as The gradient theorem can be used to establish that the electrostatic potential is the amount of work per unit charge required to move a charge from point

5406-403: The field point r {\displaystyle \mathbf {r} } , and r ^ i   = d e f   r i | r i | {\textstyle {\hat {\mathbf {r} }}_{i}\ {\stackrel {\mathrm {def} }{=}}\ {\frac {\mathbf {r} _{i}}{|\mathbf {r} _{i}|}}}

5508-473: 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,

5610-407: The secondary and tertiary structures of proteins and nucleic acids . In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named the proton acceptor, whereas the one covalently bound to the hydrogen is named the proton donor. This nomenclature is recommended by the IUPAC. The hydrogen of the donor is protic and therefore can act as a Lewis acid and the acceptor

5712-420: The substrate and the enzyme, therefore the importance of these interactions is especially great in biochemistry and molecular biology , and is the basis of enzymology ). A hydrogen bond is an extreme form of dipole-dipole bonding, referring to the attraction between a hydrogen atom that is bonded to an element with high electronegativity , usually nitrogen , oxygen , or fluorine . The hydrogen bond

5814-425: The test charge q {\displaystyle q} , which situated at the point r {\displaystyle \mathbf {r} } , and ϕ ( r ) {\displaystyle \phi (\mathbf {r} )} is the electric potential that would be at r {\displaystyle \mathbf {r} } if the test charge were not present. If only two charges are present,

5916-448: The acidic proton in the enol tautomer of acetylacetone appears at ⁠ δ H {\displaystyle \delta _{\text{H}}} ⁠  15.5, which is about 10 ppm downfield of a conventional alcohol. In the IR spectrum, hydrogen bonding shifts the X−H stretching frequency to lower energy (i.e. the vibration frequency decreases). This shift reflects

6018-535: The assumption that the electric field is irrotational : From Faraday's law , this assumption implies the absence or near-absence of time-varying magnetic fields: In other words, electrostatics does not require the absence of magnetic fields or electric currents. Rather, if magnetic fields or electric currents do exist, they must not change with time, or in the worst-case, they must change with time only very slowly . In some problems, both electrostatics and magnetostatics may be required for accurate predictions, but

6120-419: The attraction between permanent dipoles (dipolar molecules) and are temperature dependent. They consist of attractive interactions between dipoles that are ensemble averaged over different rotational orientations of the dipoles. It is assumed that the molecules are constantly rotating and never get locked into place. This is a good assumption, but at some point molecules do get locked into place. The energy of

6222-406: The attractive force. In contrast, the influence of the repulsive force is essentially unaffected by temperature. When a gas is compressed to increase its density, the influence of the attractive force increases. If the gas is made sufficiently dense, the attractions can become large enough to overcome the tendency of thermal motion to cause the molecules to disperse. Then the gas can condense to form

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6324-551: The average number of hydrogen bonds increases to 3.69. Another study found a much smaller number of hydrogen bonds: 2.357 at 25 °C. Defining and counting the hydrogen bonds is not straightforward however. 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

6426-513: The backbone amide C=O of residue i as the H-bond acceptor and two H-bond donors from residue i + 4 : the backbone amide N−H and a side-chain hydroxyl or thiol H . The energy preference of the bifurcated H-bond hydroxyl or thiol system is -3.4 kcal/mol or -2.6 kcal/mol, respectively. This type of bifurcated H-bond provides an intrahelical H-bonding partner for polar side-chains, such as serine , threonine , and cysteine within

6528-419: The basic structure of the polymer backbone. This hierarchy of bond strengths (covalent bonds being stronger than hydrogen-bonds being stronger than van der Waals forces) is relevant in the properties of many materials. 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,

6630-507: The bond length is very adaptable to the metal complex/hydrogen donor system. The Hydrogen bond is relevant to drug design. According to Lipinski's rule of five the majority of orally active drugs have no more than five hydrogen bond donors and fewer than ten hydrogen bond acceptors. These interactions exist between nitrogen – hydrogen and oxygen –hydrogen centers. Many drugs do not, however, obey these "rules". Intermolecular An intermolecular force ( IMF ; also secondary force )

6732-416: The chains. Prominent examples include cellulose and its derived fibers, such as cotton and flax . In nylon , hydrogen bonds between carbonyl and the amide N H effectively link adjacent chains, which gives the material mechanical strength. Hydrogen bonds also affect the aramid fibre , where hydrogen bonds stabilize the linear chains laterally. The chain axes are aligned along the fibre axis, making

6834-492: The charge of any ion is much greater than the charge of a dipole moment. Ion–dipole bonding is stronger than hydrogen bonding. An ion–dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, allowing maximum attraction. An important example of this interaction is hydration of ions in water which give rise to hydration enthalpy . The polar water molecules surround themselves around ions in water and

6936-417: 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

7038-496: The cohesion of condensed phases and physical absorption of gases, but also to a universal force of attraction between macroscopic bodies. The first contribution to van der Waals forces is due to electrostatic interactions between rotating permanent dipoles, quadrupoles (all molecules with symmetry lower than cubic), and multipoles. It is termed the Keesom interaction , named after Willem Hendrik Keesom . These forces originate from

7140-473: The coupling between the two can still be ignored. Electrostatics and magnetostatics can both be seen as non-relativistic Galilean limits for electromagnetism. In addition, conventional electrostatics ignore quantum effects which have to be added for a complete description. As the electric field is irrotational , it is possible to express the electric field as the gradient of a scalar function, ϕ {\displaystyle \phi } , called

7242-432: The damage of electronic components during manufacturing, and photocopier and laser printer operation. The electrostatic model accurately predicts electrical phenomena in "classical" cases where the velocities are low and the system is macroscopic so no quantum effects are involved. It also plays a role in quantum mechanics, where additional terms also need to be included. Coulomb's law states that: The magnitude of

7344-448: The dehydration stabilizes the hydrogen bond by destabilizing the nonbonded state consisting of dehydrated isolated charges . 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. The properties of many polymers are affected by hydrogen bonds within and/or between

7446-411: The donors and acceptors for hydrogen bonds on those solutes. Hydrogen bonds between water molecules have an average lifetime of 10 seconds, or 10 picoseconds. A single hydrogen atom can participate in two hydrogen bonds. This type of bonding is called "bifurcated" (split in two or "two-forked"). It can exist, for instance, in complex organic molecules. It has been suggested that a bifurcated hydrogen atom

7548-478: 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 . 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

7650-408: The electric field at r {\displaystyle \mathbf {r} } (called the field point ) of: where r i = r − r i {\textstyle \mathbf {r} _{i}=\mathbf {r} -\mathbf {r} _{i}} is the displacement vector from a source point r i {\displaystyle \mathbf {r} _{i}} to

7752-550: The electric field. Field lines begin on positive charge and terminate on negative charge. They are parallel to the direction of the electric field at each point, and the density of these field lines is a measure of the magnitude of the electric field at any given point. A collection of n {\displaystyle n} particles of charge q i {\displaystyle q_{i}} , located at points r i {\displaystyle \mathbf {r} _{i}} (called source points ) generates

7854-472: The electron density of the system. Interpretations of the anisotropies in the Compton profile of ordinary ice claim that the hydrogen bond is partly covalent. However, this interpretation was challenged and subsequently clarified. 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

7956-411: The electrostatic force of attraction or repulsion between two point charges is directly proportional to the product of the magnitudes of charges and inversely proportional to the square of the distance between them. The force is along the straight line joining them. If the two charges have the same sign, the electrostatic force between them is repulsive; if they have different signs, the force between them

8058-517: The energy released during the process is known as hydration enthalpy. The interaction has its immense importance in justifying the stability of various ions (like Cu ) in water. An ion–induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole–induced dipole force, the charge of the ion causes distortion of the electron cloud on the non-polar molecule. The van der Waals forces arise from interaction between uncharged atoms or molecules, leading not only to such phenomena as

8160-422: The fibres extremely stiff and strong. Hydrogen-bond networks make both 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 . A symmetric hydrogen bond is a special type of hydrogen bond in which

8262-559: The following types: Information on intermolecular forces is obtained by macroscopic measurements of properties like viscosity , pressure, volume, temperature (PVT) data. The link to microscopic aspects is given by virial coefficients and intermolecular pair potentials , such as the Mie potential , Buckingham potential or Lennard-Jones potential . In the broadest sense, it can be understood as such interactions between any particles ( molecules , atoms , ions and molecular ions ) in which

8364-406: The force (and hence the field) can be calculated by summing over all the contributions due to individual source particles is an example of the superposition principle . The electric field produced by a distribution of charges is given by the volume charge density ρ ( r ) {\displaystyle \rho (\mathbf {r} )} and can be obtained by converting this sum into

8466-412: The formation of chemical (that is, ionic, covalent or metallic) bonds does not occur. In other words, these interactions are significantly weaker than covalent ones and do not lead to a significant restructuring of the electronic structure of the interacting particles. (This is only partially true. For example, all enzymatic and catalytic reactions begin with a weak intermolecular interaction between

8568-438: 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

8670-581: The hydrophobic membrane environments. 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, circular dichroism (CD) experiments have shown osmolyte to act through an enthalpic effect. The molecular mechanism for their role in protein stabilization

8772-660: The identification of hydrogen bonds also in complicated molecules is crystallography , sometimes also NMR-spectroscopy. Structural details, in particular distances between donor and acceptor which are smaller than the sum of the van der Waals radii can be taken as indication of the hydrogen bond strength. One scheme gives the following somewhat arbitrary classification: those that are 15 to 40 kcal/mol, 5 to 15 kcal/mol, and >0 to 5 kcal/mol are considered strong, moderate, and weak, respectively. Hydrogen bonds involving C-H bonds are both very rare and weak. The resonance assisted hydrogen bond (commonly abbreviated as RAHB)

8874-411: The induced dipole is free to shift and rotate around the polar molecule. The Debye induction effects and Keesom orientation effects are termed polar interactions. The induced dipole forces appear from the induction (also termed polarization ), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) 31 on another. This interaction

8976-405: The ions; in contrast to many other noncovalent interactions, salt bridges are not directional and show in the solid state usually contact determined only by the van der Waals radii of the ions. Inorganic as well as organic ions display in water at moderate ionic strength I similar salt bridge as association ΔG values around 5 to 6 kJ/mol for a 1:1 combination of anion and cation, almost independent of

9078-631: The nature (size, polarizability, etc.) of the ions. The ΔG values are additive and approximately a linear function of the charges, the interaction of e.g. a doubly charged phosphate anion with a single charged ammonium cation accounts for about 2x5 = 10 kJ/mol. The ΔG values depend on the ionic strength I of the solution, as described by the Debye-Hückel equation, at zero ionic strength one observes ΔG = 8 kJ/mol. Dipole–dipole interactions (or Keesom interactions) are electrostatic interactions between molecules which have permanent dipoles. This interaction

9180-483: The onset of orientational or rotational disorder of the ions. Hydrogen bonding is of persistent 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. Quantum chemical calculations of the relevant interresidue potential constants ( compliance constants ) revealed large differences between individual H bonds of

9282-453: 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 proteins. Compared to the C−C , C−O , and C−N bonds that comprise most polymers, hydrogen bonds are far weaker, perhaps 5%. Thus, hydrogen bonds can be broken by chemical or mechanical means while retaining

9384-460: The points r i {\displaystyle \mathbf {r} _{i}} . This potential energy (in Joules ) is: where R i = r − r i {\displaystyle \mathbf {\mathcal {R_{i}}} =\mathbf {r} -\mathbf {r} _{i}} is the distance of each charge Q i {\displaystyle Q_{i}} from

9486-497: The potential energy is Q 1 Q 2 / ( 4 π ε 0 r ) {\displaystyle Q_{1}Q_{2}/(4\pi \varepsilon _{0}r)} . The total electric potential energy due a collection of N charges is calculating by assembling these particles one at a time : where the following sum from, j = 1 to N , excludes i = j : This electric potential, ϕ i {\displaystyle \phi _{i}}

9588-432: The prescription ∑ ( ⋯ ) → ∫ ( ⋯ ) ρ d 3 r {\textstyle \sum (\cdots )\rightarrow \int (\cdots )\rho \,\mathrm {d} ^{3}r} : This second expression for electrostatic energy uses the fact that the electric field is the negative gradient of the electric potential, as well as vector calculus identities in

9690-484: 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

9792-495: 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. Theoretically, the bond strength of the hydrogen bonds can be assessed using NCI index, non-covalent interactions index , which allows a visualization of these non-covalent interactions , as its name indicates, using

9894-421: The structure of polymers , both synthetic and natural. The attraction between cationic and anionic sites is a noncovalent, or intermolecular interaction which is usually referred to as ion pairing or salt bridge. It is essentially due to electrostatic forces, although in aqueous medium the association is driven by entropy and often even endothermic. Most salts form crystals with characteristic distances between

9996-427: The tetrameric quaternary structure within the mammalian sorbitol dehydrogenase protein family. 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,

10098-578: The theory of van der Waals between macroscopic bodies in 1937 and showed that the additivity of these interactions renders them considerably more long-range. (kJ/mol) This comparison is approximate. The actual relative strengths will vary depending on the molecules involved. For instance, the presence of water creates competing interactions that greatly weaken the strength of both ionic and hydrogen bonds. We may consider that for static systems, Ionic bonding and covalent bonding will always be stronger than intermolecular forces in any given substance. But it

10200-406: The total electrostatic energy only if both are integrated over all space. On a conductor , a surface charge will experience a force in the presence of an electric field . This force is the average of the discontinuous electric field at the surface charge. This average in terms of the field just outside the surface amounts to: This pressure tends to draw the conductor into the field, regardless of

10302-442: The work of a fellow scientist at their laboratory, Maurice Loyal Huggins , saying, "Mr. Huggins of this laboratory in some work as yet unpublished, has used the idea of a hydrogen kernel held between two atoms as a theory in regard to certain organic compounds." An 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. The simplest case

10404-402: Was thus the source of the word electricity . Electrostatic phenomena arise from the forces that electric charges exert on each other. Such forces are described by Coulomb's law . There are many examples of electrostatic phenomena, from those as simple as the attraction of plastic wrap to one's hand after it is removed from a package, to the apparently spontaneous explosion of grain silos,

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