Hydrocarbon

From Wikipedia, the free encyclopedia

Ball-and-stick model of the methane molecule, CH₄. Methyne is part of a homologous series known as the alkanes, which contain single bonds only.

 

In organic chemistry, a hydrocarbon is an organic compound consisting entirely of hydrogen and carbon. Hydrocarbons are examples of group 14 hydrides. Hydrocarbons from which one hydrogen atom has been removed are functional groups called hydrocarbyls. Because carbon has 4 electrons in its outermost shell (and because each covalent bond requires a donation of 1 electron, per atom, to the bond) carbon has exactly four bonds to make, and is only stable if all 4 of these bonds are used. 

Aromatic hydrocarbons (arenes), alkanes, cycloalkanes and alkyne-based compounds are different types of hydrocarbons.

Most hydrocarbons found on Earth naturally occur in crude oil, where decomposed organic matter provides an abundance of carbon and hydrogen which, when bonded, can catenate to form seemingly limitless chains.

Types of hydrocarbons

As defined by IUPAC nomenclature of organic chemistry, the classifications for hydrocarbons are:

1. Saturated hydrocarbons are the simplest of the hydrocarbon species. They are composed entirely of single bonds and are saturated with hydrogen. The formula for acyclic saturated hydrocarbons (i.e., alkanes) is C_nH_2n+2. The most general form of saturated hydrocarbons is C_nH_2n+2(1-r), where r is the number of rings. Those with exactly one ring are the cycloalkanes. Saturated hydrocarbons are the basis of petroleum fuels and are found as either linear or branched species. Substitution reaction is their characteristics property (like chlorination reaction to form chloroform). Hydrocarbons with the same molecular formula but different structural formulae are called structural isomers. As given in the example of 3-methylhexane and its higher homologues, branched hydrocarbons can be chiral. Chiral saturated hydrocarbons constitute the side chains of biomolecules such as chlorophyll and tocopherol.
2. Unsaturated hydrocarbons have one or more double or triple bonds between carbon atoms. Those with double bond are called alkenes. Those with one double bond have the formula C_nH_2n (assuming non-cyclic structures). Those containing triple bonds are called alkyne. Those with one triple bond have the formula C_nH_2n−2.
3. Aromatic hydrocarbons, also known as arenes, are hydrocarbons that have at least one aromatic ring.

Hydrocarbons can be gases (e.g. methane and propane), liquids (e.g. hexane and benzene), waxes or low melting solids (e.g. paraffin wax and naphthalene) or polymers (e.g. polyethylene, polypropylene and polystyrene).

General properties

Because of differences in molecular structure, the empirical formula remains different between hydrocarbons; in linear or "straight-run" alkanes, alkenes and alkynes, the amount of bonded hydrogen lessens in alkenes and alkynes due to the "self-bonding" or catenation of carbon preventing entire saturation of the hydrocarbon by the formation of double or triple bonds. 

This inherent ability of hydrocarbons to bond to themselves is known as catenation, and allows hydrocarbons to form more complex molecules, such as cyclohexane, and in rarer cases, arenes such as benzene. This ability comes from the fact that the bond character between carbon atoms is entirely non-polar, in that the distribution of electrons between the two elements is somewhat even due to the same electronegativity values of the elements (~0.30), and does not result in the formation of an electrophile. 

Generally, with catenation comes the loss of the total amount of bonded hydrocarbons and an increase in the amount of energy required for bond cleavage due to strain exerted upon the molecule; in molecules such as cyclohexane, this is referred to as ring strain, and occurs due to the "destabilized" spatial electron configuration of the atom. 

In simple chemistry, as per valence bond theory, the carbon atom must follow the "4-hydrogen rule", which states that the maximum number of atoms available to bond with carbon is equal to the number of electrons that are attracted into the outer shell of carbon. In terms of shells, carbon consists of an incomplete outer shell, which comprises 4 electrons, and thus has 4 electrons available for covalent or dative bonding. 

Hydrocarbons are hydrophobic like lipids. 

Some hydrocarbons also are abundant in the solar system. Lakes of liquid methane and ethane have been found on Titan, Saturn's largest moon, confirmed by the Cassini-Huygens Mission. Hydrocarbons are also abundant in nebulae forming polycyclic aromatic hydrocarbon (PAH) compounds.

Simple hydrocarbons and their variations

Number of carbon atoms	Alkane (single bond)	Alkene (double bond)	Alkyne (triple bond)	Cycloalkane	Alkadiene
1	Methane	—	—	—	—
2	Ethane	Ethene	Ethyne	—	—
3	Propane	Propene	Propyne	Cyclopropane	Propadiene
4	Butane	Butene	Butyne	Cyclobutane	Butadiene
5	Pentane	Pentene	Pentyne	Cyclopentane	Pentadiene
6	Hexane	Hexene	Hexyne	Cyclohexane	Hexadiene
7	Heptane	Heptene	Heptyne	Cycloheptane	Heptadiene
8	Octane	Octene	Octyne	Cyclooctane	Octadiene
9	Nonane	Nonene	Nonyne	Cyclononane	Nonadiene
10	Decane	Decene	Decyne	Cyclodecane	Decadiene
11	Undecane	Undecene	Undecyne	Cycloundecane	Undecadiene
12	Dodecane	Dodecene	Dodecyne	Cyclododecane	Dodecadiene

Usage

Oil refineries are one way hydrocarbons are processed for use. Crude oil is processed in several stages to form desired hydrocarbons, used as fuel and in other products.

 

Tank wagon 33 80 7920 362-0 with hydrocarbon gas at Bahnhof Enns (2018).

 

Hydrocarbons are a primary energy source for current civilizations. The predominant use of hydrocarbons is as a combustible fuel source. In their solid form, hydrocarbons take the form of asphalt (bitumen). 

Mixtures of volatile hydrocarbons are now used in preference to the chlorofluorocarbons as a propellant for aerosol sprays, due to chlorofluorocarbons' impact on the ozone layer.

Methane (CH₄) and ethane (C₂H₆) are gaseous at ambient temperatures and cannot be readily liquefied by pressure alone. Propane (C₃H₈) is however easily liquefied, and exists in 'propane bottles' mostly as a liquid. Butane (C₄H₁₀) is so easily liquefied that it provides a safe, volatile fuel for small pocket lighters. Pentane (C₅H₁₂) is a colorless liquid at room temperature, commonly used in chemistry and industry as a powerful nearly odorless solvent of waxes and high molecular weight organic compounds, including greases. Hexane (C₆H₁₄) is also a widely used non-polar, non-aromatic solvent, as well as a significant fraction of common gasoline. The C⁶ through C¹⁰ alkanes, alkenes and isomeric cycloalkanes are the top components of gasoline, naphtha, jet fuel and specialized industrial solvent mixtures. With the progressive addition of carbon units, the simple non-ring structured hydrocarbons have higher viscosities, lubricating indices, boiling points, solidification temperatures, and deeper color. At the opposite extreme from methane lie the heavy tars that remain as the lowest fraction in a crude oil refining retort. They are collected and widely utilized as roofing compounds, pavement composition, wood preservatives (the creosote series) and as extremely high viscosity shear-resisting liquids. 

Hydrocarbon use is also prevalent in nature. Some eusocial arthropods, such as the Brazilian stingless bee Schwarziana quadripunctata, use unique hydrocarbon "scents" in order to determine kin from non-kin. The chemical hydrocarbon composition varies between age, sex, nest location, and hierarchal position.

Poisoning

Hydrocarbon poisoning such as that of benzene and petroleum usually occurs accidentally by inhalation or ingestion of these cytotoxic chemical compounds. Intravenous or subcutaneous injection of petroleum compounds with intent of suicide or abuse is an extraordinary event that can result in local damage or systemic toxicity such as tissue necrosis, abscess formation, respiratory system failure and partial damage to the kidneys, the brain and the nervous system. Moaddab and Eskandarlou report a case of chest wall necrosis and empyema resulting from attempting suicide by injection of petroleum into the pleural cavity.

Reactions

There are three main types of reactions:

• Substitution reaction
• Addition reaction
• Combustion

Substitution reactions

Substitution reactions only occur in saturated hydrocarbons (single carbon–carbon bonds). In this reaction, an alkane reacts with a chlorine molecule. One of the chlorine atoms displaces a hydrogen atom. This forms hydrochloric acid as well as the hydrocarbon with one chlorine atom.

CH₄ + Cl₂ → CH₃Cl + HCl

CH₃Cl + Cl₂ → CH₂Cl₂ + HCl

all the way to CCl₄ (carbon tetrachloride)

C₂H₆ + Cl₂ → C₂H₅Cl + HCl

C₂H₄Cl₂ + Cl₂ → C₂H₃Cl₃ + HCl

all the way to C₂Cl₆ (hexachloroethane)

Addition reactions

Addition reactions involve alkenes and alkynes. In this reaction a halogen molecule breaks the double or triple bond in the hydrocarbon and forms a bond.

Combustion

Hydrocarbons are currently the main source of the world's electric energy and heat sources (such as home heating) because of the energy produced when burnt. Often this energy is used directly as heat such as in home heaters, which use either petroleum or natural gas. The hydrocarbon is burnt and the heat is used to heat water, which is then circulated. A similar principle is used to create electric energy in power plants. 

Common properties of hydrocarbons are the facts that they produce steam, carbon dioxide and heat during combustion and that oxygen is required for combustion to take place. The simplest hydrocarbon, methane, burns as follows:

CH₄ + 2 O₂ → 2 H₂O + CO₂ + energy

In inadequate supply of air, carbon monoxide gas and water vapor are formed:

2 CH₄ + 3 O₂ → 2 CO + 4 H₂O

Another example is the combustion of propane:

C₃H₈ + 5 O₂ → 4 H₂O + 3 CO₂ + energy

And finally, for any linear alkane of n carbon atoms,

C_nH_2n+2 + 3n + 1/2 O₂ → (n + 1) H₂O + n CO₂ + energy.

Burning of hydrocarbons is an example of an exothermic chemical reaction. 

Hydrocarbons can also be burned with elemental fluorine, resulting in carbon tetrafluoride and hydrogen fluoride products.

Petroleum/Petrol

Natural oil spring in Korňa, Slovakia

 

Extracted hydrocarbons in a liquid form are referred to as petroleum (literally "rock oil") or mineral oil, whereas hydrocarbons in a gaseous form are referred to as natural gas. Petroleum and natural gas are found in the Earth's subsurface with the tools of petroleum geology and are a significant source of fuel and raw materials for the production of organic chemicals.

The extraction of liquid hydrocarbon fuel from sedimentary basins is integral to modern energy development. Hydrocarbons are mined from oil sands and oil shale, and potentially extracted from sedimentary methane hydrates. These reserves require distillation and upgrading to produce synthetic crude and petroleum. 

Oil reserves in sedimentary rocks are the source of hydrocarbons for the energy, transport and petrochemical industries. 

Economically important hydrocarbons include fossil fuels such as coal, petroleum and natural gas, and its derivatives such as plastics, paraffin, waxes, solvents and oils. Hydrocarbons – along with NO_x and sunlight – contribute to the formation of tropospheric ozone and greenhouse gases.

Bioremediation

Bacteria in the gabbroic layer of the ocean's crust can degrade hydrocarbons; but the extreme environment makes research difficult. Other bacteria such as Lutibacterium anuloederans can also degrade hydrocarbons. Mycoremediation or breaking down of hydrocarbon by mycelium and mushroom is possible.

Safety

Many hydrocarbons are highly flammable, therefore, care should be taken to prevent injury. Benzene and many aromatic compounds are possible carcinogens, and proper safety equipment must be worn to prevent these harmful compounds from entering the body. If hydrocarbons undergo combustion in tight areas, toxic carbon monoxide can form. Hydrocarbons should be kept away from fluorine compounds due to the high probability of forming toxic hydrofluoric acid.

Environmental impact

Hydrocarbons are introduced into the environment through their extensive use as fuels and chemicals as well as through leaks or accidental spills during exploration, production, refining, or transport. Anthropogenic hydrocarbon contamination of soil is a serious global issue due to contaminant persistence and the negative impact on human health.

Hydrogen bond (updated)

From Wikipedia, the free encyclopedia

Model of hydrogen bonds (1) between molecules of water

 

AFM image of napthalenetetracarboxylic diimide molecules on silver-terminated silicon, interacting via hydrogen bonding, taken at 77  K. ("Hydrogen bonds" in the top image are exaggerated by artifacts of the imaging technique.)

A hydrogen bond is a partially electrostatic attraction between a hydrogen (H) atom which is bound to a more electronegative atom or group, such as nitrogen (N), oxygen (O), or fluorine (F)—the hydrogen bond donor—and another adjacent atom bearing a lone pair of electrons—the hydrogen bond acceptor. 

Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule). Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, and environment, the energy of a hydrogen bond can vary between 1 and 40 kcal/mol. This makes them somewhat stronger than a van der Waals interaction, and weaker than fully covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins. 

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have much weaker hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.

In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. This definition specifies:

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

Bonding

An example of intermolecular hydrogen bonding in a self-assembled dimer complex. The hydrogen bonds are represented by dotted lines.

 

Intramolecular hydrogen bonding in acetylacetone helps stabilize the enol tautomer.

Definitions and general characteristics

A hydrogen atom attached to a relatively electronegative atom is the hydrogen bond donor. C-H bonds only participate in hydrogen bonding when the carbon atom is bound to electronegative substituents, as is the case in chloroform, CHCl₃. 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.

Examples of hydrogen bond donating (donors) and hydrogen bond accepting groups (acceptors)

 

Cyclic dimer of acetic acid; dashed green lines represent hydrogen bonds

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.

Bond strength

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). Typical enthalpies in vapor include:

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

Structural details

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

Acceptor···donor	VSEPR geometry	Angle (°)
HCN···HF	linear	180
H2CO···HF	trigonal planar	120
H2O···HF	pyramidal	46
H2S···HF	pyramidal	89
SO2···HF	trigonal	142

Spectroscopy

Strong hydrogen bonds are revealed by downfield shifts in the ¹H NMR spectrum. For example, the acidic proton in the enol tautomer of acetylacetone appears at δ15.5, 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 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.

Theoretical considerations

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

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

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.

From interpretations of the anisotropies in the Compton profile of ordinary ice that the hydrogen bond is partly covalent. However, this interpretation was challenged.

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

History

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 cite work by 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."

Hydrogen bonds in small molecules

Crystal structure of hexagonal ice. Gray dashed lines indicate hydrogen bonds

 

Structure of nickel bis(dimethylglyoximate), which features two linear hydrogen-bonds.

Water

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. 

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.

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. 

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 (H₃O⁺). (Hydronium ions are also known as "hydroxonium" ions.)

H−O⁻ H₃O⁺

Indeed, in pure water under conditions of standard temperature and pressure, this latter formulation is applicable only rarely; on average about one in every 5.5 × 10⁸ molecules gives up a proton to another water molecule, in accordance with the value of the dissociation constant for water under such conditions. It is a crucial part of the uniqueness of water. 

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. Hydrogen bonds between water molecules have an average lifetime of 10⁻¹¹ seconds, or 10 picoseconds.

Bifurcated and over-coordinated hydrogen bonds in water

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.

Other liquids

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).

H−F···H−F···H−F

Further manifestations of solvent hydrogen bonding

• Increase in the melting point, boiling point, solubility, and viscosity of many compounds can be explained by the concept of hydrogen bonding.
• Negative azeotropy of mixtures of HF and water
• The fact that ice is less dense than liquid water is due to a crystal structure stabilized by hydrogen bonds.
• Dramatically higher boiling points of NH₃, H₂O, and HF compared to the heavier analogues PH₃, H₂S, and HCl, where hydrogen-bonding is absent.
• Viscosity of anhydrous phosphoric acid and of glycerol
• Dimer formation in carboxylic acids and hexamer formation in hydrogen fluoride, which occur even in the gas phase, resulting in gross deviations from the ideal gas law.
• Pentamer formation of water and alcohols in apolar solvents.

Hydrogen bonds in polymers

Hydrogen bonding plays an important role in determining the three-dimensional structures and the properties adopted by many synthetic and natural proteins.

DNA

The structure of part of a DNA double helix

 

Hydrogen bonding between guanine and cytosine, one of two types of base pairs in DNA

 

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.

Proteins

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 3₁₀ 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.

Bifurcated H-bond systems are common in alpha-helical transmembrane proteins between 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 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, 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.

Several studies have shown that hydrogen bonds play an important role for the stability between subunits in multimeric proteins. For example, a study of sorbitol dehydrogenase displayed an important hydrogen bonding network which stabilizes the tetrameric quaternary structure within the mammalian sorbitol dehydrogenase protein family.

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.

Wool, being a protein fiber, 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.

Cellulose

Hydrogen bonds are important in the structure of cellulose and derived polymers in its many different forms in nature, such as cotton and flax. 

Para-aramid structure

 

A strand of cellulose (conformation I_α), showing the hydrogen bonds (dashed) within and between cellulose molecules

Synthetic polymers

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 fiber, where hydrogen bonds stabilize the linear chains laterally. The chain axes are aligned along the fibre axis, making the fibres extremely stiff and strong.

The hydrogen-bond networks make both natural and synthetic polymers sensitive to humidity levels in the atmosphere because water molecules can diffuse into the surface and disrupt the network. Some polymers are more sensitive than others. Thus nylons are more sensitive than aramids, and nylon 6 more sensitive than nylon-11.

Symmetric hydrogen bond

A symmetric hydrogen bond is a special type of hydrogen bond in which the proton is spaced exactly halfway between two identical atoms. The strength of the bond to each of those atoms is equal. It is an example of a three-center four-electron bond. This type of bond is much stronger than a "normal" hydrogen bond. The effective bond order is 0.5, so its strength is comparable to a covalent bond. It is seen in ice at high pressure, and also in the solid phase of many anhydrous acids such as hydrofluoric acid and formic acid at high pressure. It is also seen in the bifluoride ion [F−H−F]⁻. 

Symmetric hydrogen bonds have been observed recently spectroscopically in formic acid at high pressure (greater than 1 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. Low-barrier hydrogen bonds form when the distance between two heteroatoms is very small.

Dihydrogen 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.

Dynamics probed by spectroscopic means

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.

Application to drugs

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 bonding phenomena

• Occurrence of proton tunneling during DNA replication is believed to be responsible for cell mutations.
• High water solubility of many compounds such as ammonia is explained by hydrogen bonding with water molecules.
• Deliquescence of NaOH is caused in part by reaction of OH⁻ with moisture to form hydrogen-bonded H
  ₃O⁻
  ₂ species. An analogous process happens between NaNH₂ and NH₃, and between NaF and HF.
• The presence of hydrogen bonds can cause an anomaly in the normal succession of states of matter for certain mixtures of chemical compounds as temperature increases or decreases. These compounds can be liquid until a certain temperature, then solid even as the temperature increases, and finally liquid again as the temperature rises over the "anomaly interval"
• Smart rubber utilizes hydrogen bonding as its sole means of bonding, so that it can "heal" when torn, because hydrogen bonding can occur on the fly between two surfaces of the same polymer.