Chemistry
Hydrogen Bonds
Hydrogen bonds are a type of weak chemical bond that occurs between a hydrogen atom and a highly electronegative atom, such as oxygen, nitrogen, or fluorine. These bonds play a crucial role in the structure and properties of many biological molecules, such as DNA and proteins, as well as in the physical properties of water.
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11 Key excerpts on "Hydrogen Bonds"
eBook - PDF- Alexander Gavrilyuk(Author)
- 2013(Publication Date)
- Jenny Stanford Publishing(Publisher)
These bonds can occur between molecules (intermolecularly), or within different parts of a single molecule (intramolecularly). This type of bond occurs both in inorganic molecules such as water and organic molecules such as DNA. New knowledge on the nature of the hydrogen bond made it possible to pose the question concerning its redefinition. In the following table there are two definitions of hydrogen bond: the current one adopted by the IUPAC and another proposed by Prof. Desiraju [33], which has good chances to substitute the current one. In the table there are abstracts both from the current and proposed definitions. Considering both definitions it is possible to conclude that the proposed definition is more “liberal” since it does not demand the necessity of an electrostatic interaction. Current definition Proposed definition A form of association between an electronegative atom and a hy-drogen atom attached to a second, relatively electronegative atom. It is best considered as an electro-static interaction, heightened by the small size of hydrogen, which permits proximity of the interact-ing dipoles or charges. The hydrogen bond is an attractive interaction between a hydro-gen 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. 2.4 Importance of the Hydrogen Bond In the world scientific literature there is bewildering amount of publications concerning the hydrogen bond, which currently grows at a great rate. Among numerous publications the classical monographs can be recommended [34–38]. Natural scientists from various branches are continuously interested in hydrogen bonding. 29 The hydrogen bond X-H···Y-Z is an attractive interaction in which an electropositive H atom intercedes between two electronegative species X and Y and brings them closer together.
eBook - PDFDihydrogen Bond
Principles, Experiments, and Applications
- Vladimir I. Bakhmutov(Author)
- 2008(Publication Date)
- Wiley-Interscience(Publisher)
PREFACE Among the various attractive forces holding molecules together, Hydrogen Bonds are the most effective, due to their pronounced directionality and relatively low bonding energies, which are particularly important for noncovalent supramolec- ular synthesis and crystal engineering. It is clear that intermolecular hydrogen bonding has a profound impact on the structure, stability, and stereochemistry of inorganic, organic, organometallic, and bioorganic molecules and molecular assemblies built via Hydrogen Bonds. Despite the relative weakness of Hydrogen Bonds (commonly estimated as 5 to 7 kcal/mol) due to cooperativity, they are responsible for the spontaneous formation of the three-dimensional shape of pro- teins and for the double helix of DNA and other complex molecular aggregates. In some sense, intermolecular Hydrogen Bonds act as glue in the buildup and design of molecular crystals. The main advantage of hydrogen-bonded crystals is the fact that they are weak and energetically flexible enough to allow annealing and editing. On the other hand, they are strong enough to impart stability to crystal systems. The role of hydrogen bonding is also well recognized in proton transfer reactions, where Hydrogen Bonds act as organizing interactions. Hydrogen bonding, one of the oldest fundamental concepts in chemistry, is constantly evolving, due to the appearance of new experimental and theoretical methods, including new approaches through computer chemistry. Dihydrogen bonding is the most intriguing discovery in this field. Although ideas about the interaction between two hydrogen atoms with opposite partial charges have been exploited by chemists for a long time, formulation of this interaction as a bonding between two hydrogen atoms was first suggested in 1993, at which time dihy- drogen bonds become objects of numerous theoretical and experimental studies.
eBook - PDF- Charlotte W. Pratt, Kathleen Cornely(Authors)
- 2017(Publication Date)
- Wiley(Publisher)
26 CHAPTER 2 Aqueous Chemistry water is less dense than other liquids because hydrogen bonding demands that individual molecules not just approach each other but interact with a certain orientation. This geometrical constraint also explains why ice floats; for other materials, the solid is denser than the liquid. Hydrogen Bonds are one type of electrostatic force Powerful covalent bonds define basic molecular constitutions, but much weaker noncovalent bonds, including Hydrogen Bonds, govern the final three-dimensional shapes of molecules and how they interact with each other. For example, about 460 kJ · mol −1 (110 kcal · mol −1 ) of en- ergy is required to break a covalent OH bond. But a hydrogen bond in water has a strength of only about 20 kJ · mol −1 (4.8 kcal · mol −1 ). Other noncovalent interactions are weaker still. Among the noncovalent interactions that occur in biological molecules are electrostatic interactions between charged groups such as carboxylate (COO − ) and amino (NH 3 + ) groups. These ionic interactions are intermediate in strength to covalent bonds and Hydrogen Bonds (Fig. 2.4). Hydrogen Bonds, despite their partial covalent nature, are classified as a type of electrostatic interaction. At about 1.8 Å, they are longer and hence weaker than a covalent OH bond (which is about 1 Å long). However, a completely noninteracting O and H would approach no closer than about 2.7 Å, which is the sum of their van der Waals radii (the van der Waals radius of an isolated atom is the distance from its nucleus to its effective electronic surface). 1 Å O H Covalent bond 2.7 Å O H No bond Hydrogen bond 1.8 Å H O (a) (b) (c) Hydrogen Bonds usually involve NH and OH groups as hydrogen donors and the electronegative N and O and occasionally S atoms as hydrogen acceptors (electronegativity is a measure of an atom’s affinity for electrons; Table 2.1 ).
- Nobuhiko Yui(Author)
- 2002(Publication Date)
- CRC Press(Publisher)
This means that the hydrogen-bonded molecules have dynamic characters around the thermodynamic equilibrium state. The weakness of the individual hydrogen bond is such that the bond is often not suf-ficient to provide the strength and specificity necessary for biological processes; however, this can be overcome by cooperation among and geometrical arrangement of a number of Hydrogen Bonds. The assembly systems of cooperative multiple hy-drogen bonds may be stable and specific enough, but small energy supplies can make them dissociate as seen in the process of replication of DNA. Such dynamism and flexibility are the most essential characteristics for sustaining life. The basic principles of the nature of Hydrogen Bonds are well described in the lit-erature. 1-6 This chapter describes the basic and essential concepts of hydrogen bond-ing. Typical examples of natural and nonnatural hydrogen-bonded supramolecular architectures are introduced along with the methodologies to construct biomimetic and biologically functional supramolecular architectures using hydrogen bonding. For more examples, see References 7 through 16. 3.2 BASIC CONCEPT OF HYDROGEN BONDING 3.2.1 D EFINITION Hydrogen Bonds are formed when a donor (D) with an available acidic hydrogen atom is brought into intimate contact with an acceptor (A) carrying an available non-bonding lone pair (Figure 3.1). In the broadest sense, the definition proposed by Pimentel and McClellan 2 is that a hydrogen bond exists when there is evidence of a bond, and evidence that this bond specifically involves a hydrogen atom already bonded to another atom. This definition does not specify the natures of the donor and acceptor atoms. Generally both D and A are highly electronegative atoms (O, N, F, Cl, Br, S) to form normal hydrogen bonding. However, the C of the hydrocarbon and the π -system may serve as a donor and an acceptor, respectively, to form weak hydrogen bonding.
eBook - PDFWater
The Forgotten Biological Molecule
- Denis Le Bihan, Hidenao Fukuyama, Denis Le Bihan, Hidenao Fukuyama(Authors)
- 2016(Publication Date)
- Jenny Stanford Publishing(Publisher)
The hydrogen bond in water is part (about 90%) electrostatic and part (about 10%) covalent. On forming the hydrogen bond, the donor hydrogen atom stretches away from its covalently bonded oxygen atom and the acceptor lone pair stretches away from its oxygen atom and toward the donor hydrogen atom, both oxygen atoms being pulled toward each other. The hydrogen-bonded proton has lower electron density relative to the other protons. Note that, even at temperatures as low as a few kelvin, there are considerable oscillations (
eBook - PDFThe Hydrogen Bond and the Water Molecule
The Physics and Chemistry of Water, Aqueous and Bio-Media
- Yves Marechal(Author)
- 2006(Publication Date)
- Elsevier Science(Publisher)
Part I THE HYDROGEN BOND This page intentionally left blank – 1 – The Hydrogen Bond: Formation, Thermodynamic Properties, Classification CHEMICAL BONDS Electrical forces that act on positively charged nuclei of various atoms and negatively charged electronic clouds that extend around these nuclei rule chemistry. The three other fundamental forces in physics, namely strong and weak interactions that act on the protons and neutrons of the nuclei and gravity, do not play any role in chemistry. The first two are much stronger than electromagnetic forces and consequently correspond to much larger energy level separations than energies due to electromagnetic interactions. It implies that in chemistry all nuclear levels are ground state levels or, in other words, nuclei are always in their fundamental state. The third fundamental force, gravity, is orders of mag-nitude too weak to have any detectable influence on electromagnetic levels. The ele-mentary constituents in chemistry are therefore atoms, made of positively charged nuclei that are always in their ground nuclear state and surrounded by negatively charged elec-tronic clouds. The precise knowledge of the structures of these electronic clouds is the object of chemistry. Atoms are the simplest arrangement of all these electrons and nuclei. They are not the most stable ones. Two H-atoms, for instance, the simplest atoms made of single protons surrounded by single electrons, are attracted to each other in such a way that their initially separated electronic clouds mix together so as to form a single cloud occupied by both electrons with different spins, which keep the two protons sepa-rated by a well-defined distance. This configuration, the H 2 molecule, is more stable by 4.5 eV than the configuration defined by the two far-away noninteracting H-atoms.
eBook - ePubUnderstanding Hydrogen Bonds
Theoretical and Experimental Views
- S?awomir J Grabowski, Sławomir J Grabowski(Authors)
- 2020(Publication Date)
- Royal Society of Chemistry(Publisher)
35One can see that various definitions and characteristics of the hydrogen bond are based on different types of parameters describing this kind of interaction. One can mention the topological criteria of the existence of the hydrogen bond proposed by Koch and Popelier;38 , 39 the definition of Weinhold and Klein2 that is based on the NBO approach,7 , 14 the spectroscopic,3 , 4 , 29 geometric,3 – 6 , 29 energetic3 – 6 and other criteria and definitions may be mentioned; some of them are described in the following chapters.It is interesting to discuss here briefly the recent definition of the hydrogen bond40 , 41 that is recommended by IUPAC and which states: ‘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’.41 As there is a large number of authors for this definition (‘Task Group on Categorizing Hydrogen Bonding and Other Intermolecular Interactions’), it is therefore called the IUPAC definition throughout this book. One can see that this definition differs only slightly from that proposed much earlier by Pimentel and McClellan.29 The hydrogen bond is labelled here as X–H Y instead of the A–H B designation; the additional matter is that ‘X is more electronegative than H’; however, this statement was included earlier in the above-cited definition of Jeffrey and Saenger.37
eBook - ePubThe Chemical Bond
Chemical Bonding Across the Periodic Table
- Gernot Frenking, Sason Shaik(Authors)
- 2014(Publication Date)
- Wiley-VCH(Publisher)
The Nature of the Chemical Bond) [8], Moore and Winmill made the first mention of the H bond, and Latimer and Rodebush recognized the importance of the H bond. Since those early days, there has been an overwhelming body of research, using X-ray crystallography, spectroscopic methods, and emerging quantum chemistry, that has made significant contributions to the enhancement of our knowledge and understanding of the H bond. However, past studies have also shown that there are many different types of H bonds and that the attributes of H bonding are much more complex and broader than previously assumed. As such, even with its history of around 100 years, the H bond is still a growing research area. The purpose of this chapter is not to discuss the diverse field of H bonding research in a comprehensive or exhaustive manner. Rather, we intend to summarize the nuts and bolts of H bonding, with some emphasis placed on the contributions of theoretical chemistry.17.2 Fundamental Properties of Hydrogen Bonds
Despite its ubiquity, it is not necessarily straightforward to define the H bond without ambiguity. A H bond can be generally described as17.1where X and Y are normally highly electronegative atoms. Although the dotted line in Eq. (17.1 ) constitutes the main part of the H bond, the H bond actually refers to the whole structural moiety composed of the three atoms, X, H, and Y [1e]. Chapter 12 of Pauling's The Nature of the Chemical Bond (third ed.) [8b] opens with the following sentences: “It was recognized some decades ago that under certain conditions an atom of hydrogen is attracted by rather strong forces to two atoms, instead of only one, so that it may be considered to be acting as a bond between them. This is called the hydrogen bond.” According to the definition by Pimentel and McClellan, a H bond exists between X–H and Y when “(a) there is evidence of bond formation (association or chelation) and (b) there is evidence that this new bond linking X–H and Y specifically involves the hydrogen atom already bonded to X [1a]”. More recently, Steiner proposed the following definition: an X–H Y interaction is called a hydrogen bond
eBook - PDF- Gabor L. Hornyak, Joydeep Dutta, H.F. Tibbals, Anil Rao(Authors)
- 2008(Publication Date)
- CRC Press(Publisher)
Physical properties like boiling point, melting point, and surface tension are influenced by Hydrogen Bonds. For example, the boiling point of Group V, VI, and VII hydrides is anomalously higher for NH 3 , H 2 O, and HF than their heavier counterparts. Table 10.8 lists some of the relevant features of Hydrogen Bonds and those of other compounds that make hydrogen bond-like bonds. 10.2.1 Standard Hydrogen Bonds Although formed from dipole–dipole interactions, the bond energy of H-bonds can be much stronger than those of standard dipole–dipole bonds: ~120 versus ~50 kJ·mol –1 , respectively. Hydrogen bond length ranges from 2.5 to 3.5 Å. This is somewhat more than the length of a normal covalent bond between hydro-gen and other electronegative atoms (<2.0 Å). H-bond length depends on the electronegative constituents of the molecules involved. For example, an H-bond between water and chlorine, a larger (softer) and therefore less electronegative atom, is generally longer and weaker than those between water and fluorine [9]. The strength of Hydrogen Bonds also depends on the immediate environment: the medium (type of liquid or gas phase), chemical composition, p H, and tem-perature. The hydrogen bond is responsible for the physical properties of water, the binding in DNA, and protein folding to form 2 ° structures. Even though the hydrogen bond is relatively weak compared to covalent bonds, the summation of all Hydrogen Bonds over the volume and surface of a system results in compounds or phases that are quite stable. In the liquid state, each water molecule is involved in an average of 3–3.5 Hydrogen Bonds—even as Hydrogen Bonds in water are formed and broken every 10 –12 s. Molecules of water pack together to form short-lived hydrogen-bonded tetra-hedral lattice transients called flickering clusters [8H 2 O ↔ 2(H 2 O) 4 ] caused by fluctuations in thermal energy. In addition to hydrogen bonding interactions, water molecules experience dispersion and polarization.
eBook - PDFHydrogen Bonding
Papers Presented at the Symposium on Hydrogen Bonding Held at Ljubljana, 29 July–3 August 1957
- D. Hadži(Author)
- 2013(Publication Date)
- Pergamon(Publisher)
I consider that there are five general conditions which are to be met in hydrogen bonding, and only there: (1) The presence of a hydrogen, or of course a deuterium or tritium atom, on or near a line joining two other atoms. According to this definition, the so-called bifurcated hydrogen bond, will be excluded and treated as an example simply of dipole attraction. (2) The atoms involved are almost exclusively those of fluorine, oxygen and nitrogen. Although it may be argued that Hydrogen Bonds are also formed with chlorine or sulphur, they are clearly of a different and generally weaker character and not enough is known about them from the point of view of crystallography to discuss them usefully here. (3) The strength of the hydrogen bond is determined by another atom, usually covalently or ionically linked to the oxygen (nitrogen, fluorine) which we may call the proton-activator. The polarizing power, in the Goldschmidt sense, of this atom determines the distance of the proton from the centre of its parent atom, the degree of screening and consequently the strength and length of the bond formed. In other terms it fixes the degree of derealization or covalent character of the hydrogen bond. Nearly all Hydrogen Bonds are subject to this effect, for in any extensively hydrogen-bonded substance, such as ice, the protons in one molecule are activated by those of its neighbours. For any given atom such as oxygen the degree of activation is greatest for the highest polarizing power. Small highly charged ions give rise to the acid-type of short hydrogen bond, while the large weak positive ions, such as the alkali metals and the tetra alkyl ammonium ions give long Hydrogen Bonds and alkaline properties. Acidity and alkalinity are reckoned relative to water, where the activation of one molecule is by means of the Hydrogen Bonds from another. The classical 2-76 A length of the hydrogen bond of ice is a kind of mean, rather than a mini-mum or maximum, hydrogen bond.
eBook - PDF- Dennis Lamb, Johannes Verlinde(Authors)
- 2011(Publication Date)
- Cambridge University Press(Publisher)
In order to account for the macroscopic properties of water, those that are measured in experiments, we must understand the nature of the bonding between molecules. Recall that the covalent bond holding each individual water molecule together is polar, so the molecule as a whole is also polar; H 2 O has a permanent dipole moment. The electroneg- ative nature of the oxygen atom concentrates the negative charge near itself, leaving the hydrogen ends of the molecule positive. Thus, the positive (H) end of each bond- ing orbital is easily attracted to one of the negatively charged lone-pair orbitals of a neighboring water molecule, as shown in Fig. 2.15. The attraction of two electronega- tive atoms (here oxygen) mediated by an intervening hydrogen atom (with partial positive charge) is termed a hydrogen bond (or H bond). When fully formed (as in ice), the H bond holds the centers of two adjacent water molecules an optimal distance of 2.76 Å (0.276 nm). About 1 Å of that separation is due to the polar covalent bond within one of the molecules, the remaining distance is accounted for by the electrostatic attraction between the H + and the lone-pair orbital of the other molecule. In the case of water, the H bond is only about 5% the strength of the covalent bond, so the H bonds between molecules break more readily than do the covalent bonds within molecules. Water thus evaporates from the liquid as intact molecules. We note at this point that the maximum number of H bonds that any one water molecule can form is two. Each water molecule has four attachment points (at the two bonding orbitals and at the two lone-pair orbitals), but each bond is shared among two neighboring molecules. Two bonds must therefore 2.1 Composition 53 be broken on average for every water molecule that evaporates from the condensed phase. The fundamental force causing the H bond is electromagnetic.
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