The octet rule is a chemical rule of thumb that reflects observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. The rule is applicable to carbon, nitrogen and the halogens, but to metals such as sodium or magnesium; the valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide. The electrons shared by the two atoms in a covalent bond are counted once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, two from the oxygen itself and two from the carbon. All four of these electrons are counted in both the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity and the second a nonmetal of high electronegativity. A chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively.
The first electron affinity of chlorine is -328.8 kJ per mole of chlorine atoms. Adding a second electron to chlorine requires energy, energy that cannot be recovered by the formation of a chemical bond; the result is that chlorine will often form a compound in which it has eight electrons in its outer shell. A sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the second electron resides in the deeper second electron shell, the second ionization energy required for its removal is much larger: +4562.4 kJ per mole. Thus sodium will, in most cases, form a compound in which it has lost a single electron and have a full outer shell of eight electrons, or octet; the energy required to transfer an electron from a sodium atom to a chlorine atom is small: +495.8 − 328.8 = +167 kJ mol−1.
This energy is offset by the lattice energy of sodium chloride: −787.3 kJ mol−1. This completes the explanation of the octet rule in this case. In the late 19th century it was known that coordination compounds were formed by the combination of atoms or molecules in such a manner that the valencies of the atoms involved became satisfied. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom is 4 or 6. In 1904 Richard Abegg was one of the first to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states that resemble the modern concept of oxidation states. Abegg noted that the difference between the maximum positive and negative valences of an element under his model is eight. In 1916, Gilbert N. Lewis referred to this insight as Abegg's rule and used it to help formulate his cubical atom model and the "rule of eight", which began to distinguish between valence and valence electrons.
In 1919 Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory". The "octet theory" evolved into what is now known as the "octet rule". Walther Kossel and Gilbert N. Lewis saw that noble gases did not have the tendency of taking part in chemical reactions under ordinary conditions. On the basis of this observation they concluded that atoms of noble gases are stable and on the basis of this conclusion they proposed a theory of valency known as "Electronic Theory of valency" in 1916: During the formation of a chemical bond, atoms combine together by gaining, losing or sharing electrons in such a way that they acquire nearest noble gas configuration; the quantum theory of the atom explains the eight electrons as a closed shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full and higher energy levels are empty. For example, the neon atom ground state has an empty n = 3 shell. According to the octet rule, the atoms before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing, or sharing electrons.
The argon atom has an analogous 3s2 3p6 configuration. There is an empty 3d level, but it is at higher energy than 3s and 3p, so that 3s2 3p6 is still considered a closed shell for chemical purposes; the atoms before and after argon tend to attain this configuration in compounds. There are, some hypervalent molecules in which the 3d level may play a part in the bonding, although this is controversial. For helium there is no 1p level according to the quantum theory, so that 1s2 is a closed shell with no p electrons; the atoms before and after helium follow a duet rule and tend to have the same 1s2 configuration as helium. The octet rule is only applicable to main group elements and there are many molecules that do not obey the octet rule; these molecules can be divided into two types: unstable intermediates that react so as to attain stability, stable molecules that follow other electron counting rules. Although stable odd-electron molecules and hypervalent molecules are taught as v
International Standard Serial Number
An International Standard Serial Number is an eight-digit serial number used to uniquely identify a serial publication, such as a magazine. The ISSN is helpful in distinguishing between serials with the same title. ISSN are used in ordering, interlibrary loans, other practices in connection with serial literature; the ISSN system was first drafted as an International Organization for Standardization international standard in 1971 and published as ISO 3297 in 1975. ISO subcommittee TC 46/SC 9 is responsible for maintaining the standard; when a serial with the same content is published in more than one media type, a different ISSN is assigned to each media type. For example, many serials are published both in electronic media; the ISSN system refers to these types as electronic ISSN, respectively. Conversely, as defined in ISO 3297:2007, every serial in the ISSN system is assigned a linking ISSN the same as the ISSN assigned to the serial in its first published medium, which links together all ISSNs assigned to the serial in every medium.
The format of the ISSN is an eight digit code, divided by a hyphen into two four-digit numbers. As an integer number, it can be represented by the first seven digits; the last code digit, which may be 0-9 or an X, is a check digit. Formally, the general form of the ISSN code can be expressed as follows: NNNN-NNNC where N is in the set, a digit character, C is in; the ISSN of the journal Hearing Research, for example, is 0378-5955, where the final 5 is the check digit, C=5. To calculate the check digit, the following algorithm may be used: Calculate the sum of the first seven digits of the ISSN multiplied by its position in the number, counting from the right—that is, 8, 7, 6, 5, 4, 3, 2, respectively: 0 ⋅ 8 + 3 ⋅ 7 + 7 ⋅ 6 + 8 ⋅ 5 + 5 ⋅ 4 + 9 ⋅ 3 + 5 ⋅ 2 = 0 + 21 + 42 + 40 + 20 + 27 + 10 = 160 The modulus 11 of this sum is calculated. For calculations, an upper case X in the check digit position indicates a check digit of 10. To confirm the check digit, calculate the sum of all eight digits of the ISSN multiplied by its position in the number, counting from the right.
The modulus 11 of the sum must be 0. There is an online ISSN checker. ISSN codes are assigned by a network of ISSN National Centres located at national libraries and coordinated by the ISSN International Centre based in Paris; the International Centre is an intergovernmental organization created in 1974 through an agreement between UNESCO and the French government. The International Centre maintains a database of all ISSNs assigned worldwide, the ISDS Register otherwise known as the ISSN Register. At the end of 2016, the ISSN Register contained records for 1,943,572 items. ISSN and ISBN codes are similar in concept. An ISBN might be assigned for particular issues of a serial, in addition to the ISSN code for the serial as a whole. An ISSN, unlike the ISBN code, is an anonymous identifier associated with a serial title, containing no information as to the publisher or its location. For this reason a new ISSN is assigned to a serial each time it undergoes a major title change. Since the ISSN applies to an entire serial a new identifier, the Serial Item and Contribution Identifier, was built on top of it to allow references to specific volumes, articles, or other identifiable components.
Separate ISSNs are needed for serials in different media. Thus, the print and electronic media versions of a serial need separate ISSNs. A CD-ROM version and a web version of a serial require different ISSNs since two different media are involved. However, the same ISSN can be used for different file formats of the same online serial; this "media-oriented identification" of serials made sense in the 1970s. In the 1990s and onward, with personal computers, better screens, the Web, it makes sense to consider only content, independent of media; this "content-oriented identification" of serials was a repressed demand during a decade, but no ISSN update or initiative occurred. A natural extension for ISSN, the unique-identification of the articles in the serials, was the main demand application. An alternative serials' contents model arrived with the indecs Content Model and its application, the digital object identifier, as ISSN-independent initiative, consolidated in the 2000s. Only in 2007, ISSN-L was defined in the
Irving Langmuir was an American chemist and physicist. He was awarded the Nobel Prize in Chemistry in 1932 for his work in surface chemistry. Langmuir's most famous publication is the 1919 article "The Arrangement of Electrons in Atoms and Molecules" in which, building on Gilbert N. Lewis's cubical atom theory and Walther Kossel's chemical bonding theory, he outlined his "concentric theory of atomic structure". Langmuir became embroiled in a priority dispute with Lewis over this work. While at General Electric from 1909 to 1950, Langmuir advanced several fields of physics and chemistry, invented the gas-filled incandescent lamp and the hydrogen welding technique; the Langmuir Laboratory for Atmospheric Research near Socorro, New Mexico, was named in his honor, as was the American Chemical Society journal for surface science called Langmuir. Irving Langmuir was born in Brooklyn, New York, on January 31, 1881, he was the third of the four children of Sadie, née Comings. During his childhood, Langmuir's parents encouraged him to observe nature and to keep a detailed record of his various observations.
When Irving was eleven, it was discovered. When this problem was corrected, details that had eluded him were revealed, his interest in the complications of nature was heightened. During his childhood, Langmuir was influenced by Arthur Langmuir. Arthur was a research chemist how things work. Arthur helped Irving set up his first chemistry lab in the corner of his bedroom, he was content to answer the myriad questions that Irving would pose. Langmuir's hobbies included mountaineering, piloting his own plane, classical music. In addition to his professional interest in the politics of atomic energy, he was concerned about wilderness conservation. Langmuir attended several schools and institutes in America and Paris before graduating high school from Chestnut Hill Academy, an elite private school located in the affluent Chestnut Hill area in Philadelphia, he graduated with a Bachelor of Science degree in metallurgical engineering from the Columbia University School of Mines in 1903. He earned his Ph.
D. in 1906 under Friedrich Dolezalek in Göttingen, for research done using the "Nernst glower", an electric lamp invented by Nernst. His doctoral thesis was entitled "On the Partial Recombination of Dissolved Gases During Cooling." He did postgraduate work in chemistry. Langmuir taught at Stevens Institute of Technology in Hoboken, New Jersey, until 1909, when he began working at the General Electric research laboratory, his initial contributions to science came from his study of light bulbs. His first major development was the improvement of the diffusion pump, which led to the invention of the high-vacuum rectifier and amplifier tubes. A year he and colleague Lewi Tonks discovered that the lifetime of a tungsten filament could be lengthened by filling the bulb with an inert gas, such as argon, the critical factor being the need for extreme cleanliness in all stages of the process, he discovered that twisting the filament into a tight coil improved its efficiency. These were important developments in the history of the incandescent light bulb.
His work in surface chemistry began at this point, when he discovered that molecular hydrogen introduced into a tungsten-filament bulb dissociated into atomic hydrogen and formed a layer one atom thick on the surface of the bulb. His assistant in vacuum tube research was his cousin William Comings White; as he continued to study filaments in vacuum and different gas environments, he began to study the emission of charged particles from hot filaments. He was one of the first scientists to work with plasmas, he was the first to call these ionized gases by that name because they reminded him of blood plasma. Langmuir and Tonks discovered electron density waves in plasmas that are now known as Langmuir waves, he introduced the concept of electron temperature and in 1924 invented the diagnostic method for measuring both temperature and density with an electrostatic probe, now called a Langmuir probe and used in plasma physics. The current of a biased probe tip is measured as a function of bias voltage to determine the local plasma temperature and density.
He discovered atomic hydrogen, which he put to use by inventing the atomic hydrogen welding process. Plasma welding has since been developed into gas tungsten arc welding. In 1917, he published a paper on the chemistry of oil films that became the basis for the award of the 1932 Nobel Prize in chemistry. Langmuir theorized that oils consisting of an aliphatic chain with a hydrophilic end group were oriented as a film one molecule thick upon the surface of water, with the hydrophilic group down in the water and the hydrophobic chains clumped together on the surface; the thickness of the film could be determined from the known volume and area of the oil, which allowed investigation of the molecular configuration before spectroscopic techniques were available. Following World War I Langmuir contributed to atomic theory and the understanding of atomic structure by defining the modern concept of valence shells and isotopes. Langmuir was president of the Institute of Radio Engineers in 1923. Based on his work at General Electric, John B. Taylor developed a detector ionizing beams of alkali metals, cal
Azide is the anion with the formula N−3. It is the conjugate base of hydrazoic acid. N−3 is a linear anion, isoelectronic with CO2, NCO−, N2O, NO+2 and NCF. Per valence bond theory, azide can be described by several resonance structures. Azide is a functional group in organic chemistry, RN3; the dominant application of azides is as a propellant in air bags. Sodium azide is made industrially by the reaction of nitrous oxide, N2O with sodium amide in liquid ammonia as solvent: N2O + 2 NaNH2 → NaN3 + NaOH + NH3Many inorganic azides can be prepared directly or indirectly from sodium azide. For example, lead azide, used in detonators, may be prepared from the metathesis reaction between lead nitrate and sodium azide. An alternative route is direct reaction of the metal with silver azide dissolved in liquid ammonia; some azides are produced by treating the carbonate salts with hydrazoic acid. The principal source of the azide moiety is sodium azide; as a pseudohalogen compound, sodium azide displaces an appropriate leaving group to give the azido compound.
Aryl azides may be prepared by displacement of the appropriate diazonium salt with sodium azide, or trimethylsilyl azide. Anilines and aromatic hydrazines undergo diazotization, as do alkyl hydrazines. Appropriately functionalized aliphatic compounds undergo nucleophilic substitution with sodium azide. Aliphatic alcohols give azides via a variant of the Mitsunobu reaction, with the use of hydrazoic acid. Hydrazines may form azides by reaction with sodium nitrite: PhNHNH2 → PhN3Alkyl or aryl acyl chlorides react with sodium azide in aqueous solution to give acyl azides, which give isocyanates in the Curtius rearrangement; the azo-transfer compounds, trifluoromethanesulfonyl azide and imidazole-1-sulfonyl azide, are prepared from sodium azide as well. They react with amines to give the corresponding azides: RNH2 → RN3 A classic method for the synthesis of azides is the Dutt–Wormall reaction in which a diazonium salt reacts with a sulfonamide first to a diazoaminosulfinate and on hydrolysis the azide and a sulfinic acid.
Azide salts can decompose with release of nitrogen gas. The decomposition temperatures of the alkali metal azides are: NaN3, KN3, RbN3, CsN3; this method is used to produce ultrapure alkali metals. Protonation of azide salts gives toxic hydrazoic acid in the presence of strong acids: H+ + N−3 → HN3Azide salts may react with heavy metals or heavy metal compounds to give the corresponding azides, which are more shock sensitive than sodium azide alone, they decompose with sodium nitrite. This is a method of destroying residual azides, prior to disposal. 2 NaN3 + 2 HNO2 → 3 N2 + 2 NO + 2 NaOHMany inorganic covalent azides have been described. The azide anion behaves as a nucleophile, it reacts with epoxides. Azides can be used as precursors of the metal nitrido complexes. Azide complexes thus is induced to release N2, generating a metal complex in unusual oxidation states. Organic azides engage in useful organic reactions; the terminal nitrogen is mildly nucleophilic. Azides extrude diatomic nitrogen, a tendency, exploited in many reactions such as the Staudinger ligation or the Curtius rearrangement or for example in the synthesis of γ-imino-β-enamino esters.
Azides may be reduced with a phosphine in the Staudinger reaction. This reaction allows azides to serve as protected -NH2 synthons, as illustrated by the synthesis of 1,1,1-trisethane: 3 H2 + CH3C3 → CH3C3 + 3 N2In the azide alkyne Huisgen cycloaddition, organic azides react as 1,3-dipoles, reacting with alkynes to give substituted 1,2,3-triazoles; this reaction is popular in click chemistry. Another azide regular is tosyl azide here in reaction with norbornadiene in a nitrogen insertion reaction: About 250 tons of azide-containing compounds are produced annually, the main product being sodium azide. Sodium azide is the propellant in automobile airbags, it decomposes on heating to give nitrogen gas, used to expand the air bag: 2 NaN3 → 2 Na + 3 N2Heavy metal salts, such as lead azide, Pb2, are shock-sensitive detonators which decompose to the corresponding metal and nitrogen, for example: Pb2 → Pb + 3 N2Silver and barium salts are used similarly. Some organic azides are an example being 2-dimethylaminoethylazide.
Because of the hazards associated with their use, few azides are used commercially although they exhibit interesting reactivity for researchers. Low molecular weight azides are considered hazardous and are avoided. In the research laboratory, azides are precursors to amines, they are popular for their participation in the "click reaction" and in Staudinger ligation. These two reactions are quite reliable, lending themselves to combinatorial chemistry; the antiviral drug zidovudine contains an azido group. Some azides are valuable as bioorthogonal chemical reporters. Azides are toxins. Sodium azide can be absorbed through the skin, it decomposes explosively upon heating to above 275 °C and reac
A carboxylate is the conjugate base of a carboxylic acid. Carboxylate salts have the general formula Mn, where M is a metal and n is 1, 2.... R and R′ are organic groups. A carboxylate ion is the conjugate base of a carboxylic acid, RCOO−, it is an ion with negative charge. A carboxylate can be made by deprotonation of carboxylic acids. Most of them has a pKa of 5, which means that they can be deprotonated by many bases, such as sodium hydroxide. Carboxylic acids dissociate into a carboxylate anion and a positively charged hydrogen ion, much more than alcohols do, because the carboxylate ion is stabilized by resonance; the negative charge, left after deprotonation of the carboxyl group is delocalized between the two electronegative oxygen atoms in a resonance structure. This delocalization of the electron cloud means that both of the oxygen atoms are less negatively charged. In contrast, an alkoxide ion, once formed, would have a strong negative charge on the oxygen atom, which would make it difficult for the proton to escape.
Carboxylic acids thus have a lower pKa values than alcohols. For example, the pKa value of ethanol is 16 while acetic acid has a pKa of 4.9. Hence acetic acid is a much stronger acid than ethanol; the higher the number of protons in solution, the lower the pH. Formate ion, HCOO− Acetate ion, CH3COO− Lactate ion, CH3CHCOO− Oxalate ion, 2−2 Citrate ion, C3H5O3−3
The SN2 reaction is a type of reaction mechanism, common in organic chemistry. In this mechanism, one bond is broken and one bond is formed synchronously, i.e. in one step. SN2 is a kind of nucleophilic substitution reaction mechanism. Since two reacting species are involved in the slow step, this leads to the term substitution nucleophilic or SN2, the other major kind is SN1. Many other more specialized mechanisms describe substitution reactions; the reaction type is so common that it has other names, e.g. "bimolecular nucleophilic substitution", or, among inorganic chemists, "associative substitution" or "interchange mechanism". The reaction most occurs at an aliphatic sp3 carbon center with an electronegative, stable leaving group attached to it, a halide atom; the breaking of the C–X bond and the formation of the new bond occur through a transition state in which a carbon under nucleophilic attack is pentacoordinate, sp2 hybridised. The nucleophile attacks the carbon at 180° to the leaving group, since this provides the best overlap between the nucleophile's lone pair and the C–X σ* antibonding orbital.
The leaving group is pushed off the opposite side and the product is formed with inversion of the tetrahedral geometry at the central atom. If the substrate under nucleophilic attack is chiral this leads to inversion of configuration, called a Walden inversion. In an example of the SN2 reaction, the attack of Br− on an ethyl chloride results in ethyl bromide, with chloride ejected as the leaving group.: SN2 attack occurs if the backside route of attack is not sterically hindered by substituents on the substrate. Therefore, this mechanism occurs at unhindered primary and secondary carbon centres. If there is steric crowding on the substrate near the leaving group, such as at a tertiary carbon centre, the substitution will involve an SN1 rather than an SN2 mechanism. Four factors affect the rate of the reaction: The substrate plays the most important part in determining the rate of the reaction; this is because the nucleophile attacks from the back of the substrate, thus breaking the carbon-leaving group bond and forming the carbon-nucleophile bond.
Therefore, to maximise the rate of the SN2 reaction, the back of the substrate must be as unhindered as possible. Overall, this means that methyl and primary substrates react the fastest, followed by secondary substrates. Tertiary substrates do not participate in SN2 reactions, because of steric hindrance. Structures that can form stable cations by simple loss of the leaving group, for example, as a resonance-stabilized carbocation, are likely to react via an SN1 pathway in competition with SN2. Like the substrate, steric hindrance affects the nucleophile's strength; the methoxide anion, for example, is both a strong base and nucleophile because it is a methyl nucleophile, is thus much unhindered. Tert-Butoxide, on the other hand, is a strong base, but a poor nucleophile, because of its three methyl groups hindering its approach to the carbon. Nucleophile strength is affected by charge and electronegativity: nucleophilicity increases with increasing negative charge and decreasing electronegativity.
For example, OH− is a better nucleophile than water, I− is a better nucleophile than Br−. In a polar aprotic solvent, nucleophilicity increases up a column of the periodic table as there is no hydrogen bonding between the solvent and nucleophile. I − would therefore be a weaker nucleophile than Br −. Verdict - A strong/anionic nucleophile always favours SN2 manner of nucleophillic substitution; the solvent affects the rate of reaction because solvents may or may not surround a nucleophile, thus hindering or not hindering its approach to the carbon atom. Polar aprotic solvents, like tetrahydrofuran, are better solvents for this reaction than polar protic solvents because polar protic solvents will hydrogen bond to the nucleophile, hindering it from attacking the carbon with the leaving group. A polar aprotic solvent with low dielectric constant or a hindered dipole end will favour SN2 manner of nucleophilic substitution reaction. Examples: DMSO, DMF, acetone etc. In polar aprotic solvent, nucleophilicity parallels basicity.
The stability of the leaving group as an anion and the strength of its bond to the carbon atom both affect the rate of reaction. The more stable the conjugate base of the leaving group is, the more that it will take the two electrons of its bond to carbon during the reaction. Therefore, the weaker the leaving group is as a conjugate base, thus the stronger its corresponding acid, the better the leaving group. Examples of good leaving groups are therefore the halides and tosylate, whereas HO− and H2N− are not; the rate of an SN2 reaction is second order, as the rate-determining step depends on the nucleophile concentration, as well as the concentration of substrate. R = kThis is a key difference between the SN2 mechanisms. In the SN1 reaction the nucleophile attacks after the rate-limiting step is over, whereas in SN2 the nucleophile forces off the leaving group in the limiting step. In other words, the rate of SN1 reactions depend only on the concentration of the substrate while the SN2 reaction rate depends on the concentration of both the substrate and nucleophile.
It has been shown that except in uncommon primary and secondary
A hydrogen bond is a electrostatic force of attraction between a hydrogen atom, covalently bound to a more electronegative atom or group the second-row elements nitrogen, oxygen, or fluorine —the hydrogen bond donor —and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor. Such an interacting system is denoted Dn–H···Ac, where the solid line denotes a covalent bond, the dotted line indicates the hydrogen bond. There is general agreement that there is a minor covalent component to hydrogen bonding for moderate to strong hydrogen bonds, although the importance of covalency in hydrogen bonding is debated. At the opposite end of the scale, there is no clear boundary between a weak hydrogen bond and a van der Waals interaction. Weaker hydrogen bonds are known for hydrogen atoms bound to elements such as chlorine; the hydrogen bond is responsible for many of the anomalous physical and chemical properties of compounds of N, O, F. Hydrogen bonds can be intramolecular.
Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, 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, weaker than 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 compared to the other group 16 hydrides that have much weaker hydrogen bonds. Intramolecular hydrogen bonding is responsible for the secondary and tertiary structures of proteins and nucleic acids, it plays an important role in the structure of polymers, both synthetic and natural. It was recognized that there are many examples of weaker hydrogen bonding involving donor Dn other than N, O, or F and/or acceptor Ac with close to or the same electronegativity as hydrogen. Though they are quite weak, they are ubiquitous and are recognized as important control elements in receptor-ligand interactions in medicinal chemistry or intra-/intermolecular interactions in materials sciences.
Thus, there is a trend of gradual broadening for the definition of hydrogen bonding. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, 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, an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation. Most introductory textbooks still restrict the definition of hydrogen bond to the "classical" type of hydrogen bond characterized in the opening paragraph. A hydrogen atom attached to a 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. Liquids that display hydrogen bonding are called associated liquids; the hydrogen bond is described as an electrostatic dipole-dipole interaction. However, it has some features of covalent bonding: it is directional and strong, produces interatomic distances shorter than the sum of the van der Waals radii, 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. Hydrogen bonds can vary in strength from weak to strong. Typical enthalpies in vapor include: F−H···:F, illustrated uniquely by HF2−, bifluoride O−H···:N, illustrated water-ammonia O−H···:O, illustrated water-water, alcohol-alcohol N−H···:N, illustrated by ammonia-ammonia N−H···:O, illustrated water-amide HO−H···:OH+3 The strength of intermolecular hydrogen bonds is most evaluated by measurements of equilibria between molecules containing donor and/or acceptor units, most 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 the identification of hydrogen bonds in complicated molecules is crystallography, sometimes 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, >0 to 5 kcal/mol are considered strong, moder