Physics
Hydrogen Ionic Bond
A hydrogen ionic bond is a type of chemical bond formed between hydrogen and another atom, typically a more electronegative one. In this bond, the hydrogen atom loses its electron to the other atom, resulting in the formation of a positively charged hydrogen ion and a negatively charged ion. This type of bond is important in various chemical reactions and compounds.
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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 - 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 - PDFAn Introduction to the Electron Theory of Solids
Metallurgy Division
- John Stringer, W. S. Owen, D. W. Hopkins, H. M. Finniston(Authors)
- 2013(Publication Date)
- Pergamon(Publisher)
This hydrogen bond is quite weak, having a bond energy of only 5 kcal/mol; but it is respon-sible for such important processes as the polymerisation in water and in hydrogen fluoride ; and is also important in certain ferro-electric crystals, notably potassium dihydrogen phosphate. The hydrogen bond may be qualitatively visualised as follows : if the normal two-electron bond between a hydrogen atom and another atom is sufficiently heteropolar, the hydrogen atom may be regarded as a bare proton, attracted to the negative ion and forming a stable pair with it at the characteristic bond distance. The proton may then be thought of as a point cation, the other atom as an anion sphere of finite radius. The positive charge on the proton is capable of attracting other anions, but when another has been attracted the two anion spheres are in contact, inhibiting the attraction of other anions to the proton. It follows that only molecules in which the normal hydrogen bond is noticeably heteropolar can form a hydrogen bridge. The Metallic Bond No summary of bonding between atoms would be complete without a reference to metallic bonding, but unfortunately this bond is very difficult to describe in the semi-qualitative terms 126 ELECTRON THEORY OF SOLIDS that have been used for the other types of bond listed above. The simplest concept of the metallic bond is that of the resonating covalent bonds, developed by Pauling, which explains the con-ductivity of metals and the non-directional nature of the bonding which is responsible for their high ductility. The model is, however, of little value for analytical purposes. A complete analysis is one of the more advanced problems in quantum mechanics, and is beyond the scope of the present book.
eBook - ePubUnderstanding Hydrogen Bonds
Theoretical and Experimental Views
- S?awomir J Grabowski, Sławomir J Grabowski(Authors)
- 2020(Publication Date)
- Royal Society of Chemistry(Publisher)
Chapter 7 ). Thus, one can see from several comparisons of classifications and definitions presented so far concerning the hydrogen bond that this is a complex phenomenon that is difficult to be categorized and defined.In very early studies there has already been trials to present the characteristic properties of the hydrogen bond. Lewis, who probably used the term ‘hydrogen bond’ for the first time, has pointed out that the hydrogen atom when connected with an extremely negative element such as fluorine, oxygen or nitrogen towards which the bond electron pair is shifted may be attached in such a case to another pair of electrons.15 , † The term ‘extremely negative element’ appears in the study of Lewis15 because the definition of electronegativity was introduced later by Pauling.1 Pauling also presented a few characteristics of the hydrogen bond1 as he pointed out that this interaction occurs if the H-atom is located between two electronegative atoms and that it interacts more strongly with one of them forming a covalent bond while its interaction with the second electronegative atom is much weaker. The hydrogen atom ‘may be considered to be acting as a bond’ between these electronegative atoms and this is called the hydrogen bond.1 On the other hand, in his earlier study, Lewis attributed the term hydrogen bond only to the second weaker interaction. In a further study, Desiraju suggests that the use of the term hydrogen bridge is preferable probably to the use of the term hydrogen bond.16
- C. Woodward, N.M. Allewell(Authors)
- 1998(Publication Date)
- Elsevier Science(Publisher)
The term electrostatic inter- actions usually implies interactions between ionizable groups; hydrogen bond refers to the interaction of a donor proton bearing partial positive charge with the electron density of an acceptor polar atom. Separation of these two types of noncovalent interactions is somewhat artificial as hydrogen bonds are fundamen- tally electrostatic. However, since the physical origins and the structural conse- quences of these two forces in proteins are significantly different, it is therefore convenient to distinguish between them when studying or interpreting their role in structure and function. Electrostatic interactions and hydrogen bonding play important roles in many aspects of folding, stability, and activity of proteins. There has been debate in the past about the exact magnitude of the contribution by hydrogen bonding to folding vis-a-vis the strongly stabilizing contribution by the hydrophobic effect. Currently, hydrogen bonds are thought to play a preeminent role in determining specificity as well as stability of the native states of proteins. The magnitude of contributions by electrostatic interactions to stability of folded structures, on the other hand, are generally believed to be modest compared to those by hydrophobic, van der Waals, or hydrogen bonding interactions. Electrostatics mediates interactions between macromolecules and polar and ionic components of their environment, including water, protons, and salts, and they play a fundamental role in the physiological regulation of many proteins. The molecular origins and physical character of electrostatic interactions in proteins are relatively well understood. Algorithms are even available for quantita- tion of electrostatic energies from X-ray elucidated structures based on principles of classical electrostatics and statistical mechanics.
eBook - PDF- Jochen Winkler(Author)
- 2014(Publication Date)
- Vincentz Network(Publisher)
In liquid water, clusters of approximately nine individual molecules exist. The high conductivity of water is, for example, attributed to the switching of hydrogen bonds to chemical bonds, whereby positive charges are transferred without mass transport being involved. Strong hydrogen bonds of the type X-H----Y occur mainly when the atom X is either an oxygen, a nitrogen or a halogen atom (fluorine, chlorine, bromine or iodine) whereas the atom Y may be oxygen, nitrogen, sulfur or a halogen atom. Hydrogen bonds play an important role in many processes taking place in the ani-mated and non-animated nature. In colloidal chemistry, hydrogen bonds are often determining factors as well because of their strength and their special alignment. There is as yet no coherent method for calculating interaction energies from hydrogen bonds between molecules. They usually lie in the order of about 10 to 50 KJ/mole. The F-H------F bond is the most powerful hydrogen bond. A bonding energy of 160 to 170 KJ/mole is attributed to it. 1.5 Range of physical interaction energies Polarizabilities, dipole moments and values of h · v 0 for a number of molecules as well as the contributions of attraction energies coming from the different mechanisms are presented in Table 1.4. It shows that dispersive forces always play a major role in intermolecular interactions. In contrast, appreciable dipole-dipole interactions only come into being when dipole moments exceed 1.3 Debye ( = 4.3 · 10-30 C · m ). Induced dipole interactions are very weak in all cases. The noble gases naturally have only dispersive interactions, whereas, on the other hand, water has a high contribution of dipolar interactions. Next to this, water molecules are attracted to each other by hydrogen bonds. An indication for the interaction between molecules or atoms is their boiling points. Table 1.5 lists polarizabilities, dipole moments and boiling points of a few selected chemical compounds.
- Morris Hein, Scott Pattison, Susan Arena, Leo R. Best(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
Molecules of H 2 O, HBr, and ICl are polar: H Cl H Br I Cl H O H How do we know whether a bond between two atoms is ionic or covalent? The differ- ence in electronegativity between the two atoms determines the character of the bond formed between them. As the difference in electronegativity increases, the polarity of the bond (or percent ionic character) increases. In general, if the electronegativity difference between two bonded atoms is greater than 1.7–1.9, the bond will be more ionic than covalent. If the electronegativity difference is greater than 2.0, the bond is strongly ionic. If the electro- negativity difference is less than 1.5, the bond is strongly covalent. P R A C T I C E 1 1 . 7 Explain the term electronegativity and how it relates to the elements in the periodic table. TABLE 11.5 Three-Dimensional Representation of Electronegativity 1 H 2.1 3 Li 1.0 11 Na 0.9 19 K 0.8 37 Rb 0.8 55 Cs 0.7 87 Fr 0.7 88 Ra 0.9 89–103 Ac–Lr 1.1–1.7 56 Ba 0.9 57–71 La–Lu 1.1–1.2 38 Sr 1.0 39 Y 1.2 40 Zr 1.4 41 Nb 1.6 42 Mo 1.8 43 Tc 1.9 44 Ru 2.2 46 Pd 2.2 47 Ag 1.9 48 Cd 1.7 49 In 1.7 50 Sn 1.8 51 Sb 1.9 52 Te 2.1 53 I 2.5 72 Hf 1.3 73 Ta 1.5 74 W 1.7 75 Re 1.9 76 Os 2.2 45 Rh 2.2 77 Ir 2.2 78 Pt 2.2 79 Au 2.4 80 Hg 1.9 81 Tl 1.8 82 Pb 1.8 83 Bi 1.9 84 Po 2.0 85 At 2.2 20 Ca 1.0 21 Sc 1.3 22 Ti 1.4 23 V 1.6 24 Cr 1.6 25 Mn 1.5 26 Fe 1.8 27 Co 1.8 28 Ni 1.8 29 Cu 1.9 30 Zn 1.6 31 Ga 1.6 32 Ge 1.8 33 As 2.0 34 Se 2.4 35 Br 2.8 13 Al 1.5 5 B 2.0 14 Si 1.8 6 C 2.5 15 P 2.1 16 S 2.5 17 Cl 3.0 7 N 3.0 8 O 3.5 9 F 4.0 4 Be 1.5 12 Mg 1.2 9 F 4.0 Atomic number Symbol Electronegativity H 2 Cl 2 HCl NaCl + – Polar covalent molecule Ionic compound Nonpolar molecules Figure 11.10 Nonpolar, polar covalent, and ionic compounds. 228 CHAPTER 11 • Chemical Bonds: The Formation of Compounds from Atoms Care must be taken to distinguish between polar bonds and polar molecules. A covalent bond between different kinds of atoms is always polar.
eBook - PDFFundamentals of Materials Science and Engineering
An Integrated Approach
- William D. Callister, Jr., David G. Rethwisch(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
Secondary or physical forces and energies are also found in many solid materials; they are weaker than the primary ones but nonetheless influence the physical properties of some materials. The sections that follow explain the several kinds of primary and secondary interatomic bonds. bonding energy primary bond 2.6 | | PRIMARY INTERATOMIC BONDS Ionic Bonding Ionic bonding is perhaps the easiest to describe and visualize. It is always found in compounds composed of both metallic and nonmetallic elements, elements situated at the horizontal extremities of the periodic table. Atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms. In the process, all the atoms acquire stable or inert gas configurations (i.e., completely filled orbital shells) and, in addition, an electrical charge—that is, they become ions. Sodium chloride (NaCl) is the classic ionic material. A sodium atom can assume the electron structure of neon (and a net single positive charge with a reduction in size) by a transfer of its one va- lence 3s electron to a chlorine atom (Figure 2.13a). After such a transfer, the chlorine ion acquires a net negative charge, an electron configuration identical to that of argon; ionic bonding 2.6 Primary Interatomic Bonds 35 it is also larger than the chlorine atom. Ionic bonding is illustrated schematically in Figure 2.13b. The attractive bonding forces are coulombic—that is, positive and negative ions, by virtue of their net electrical charge, attract one another. For two isolated ions, the attrac- tive energy E A is a function of the interatomic distance according to E A = − A __ r (2.9) Theoretically, the constant A is equal to A = 1 ____ 4 π ε 0 (|Z 1 |e)(|Z 2 |e) (2.10) Here ε 0 is the permittivity of a vacuum (8.85 × 10 −12 F/m), |Z 1 | and |Z 2 | are absolute values of the valences for the two ion types, and e is the electronic charge (1.602 × 10 −19 C).
eBook - PDF- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
When two different kinds of atoms share a pair of electrons, a bond forms in which electrons are shared unequally. One atom assumes a partial positive charge and the other a partial nega- tive charge with respect to each other. This difference in charge occurs because the two atoms exert unequal attraction for the pair of shared electrons. The attractive force that an atom of an element has for shared electrons in a molecule or polyatomic ion is known as its electro- negativity. Elements differ in their electronegativities. For example, both hydrogen and chlo- rine need one electron to form stable electron configurations. They share a pair of electrons in hydrogen chloride (HCl). Chlorine is more electronegative and therefore has a greater attrac- tion for the shared electrons than does hydrogen. As a result, the pair of electrons is displaced toward the chlorine atom, giving it a partial negative charge and leaving the hydrogen atom with a partial positive charge. Note that the electron is not transferred entirely to the chlorine atom (as in the case of sodium chloride) and that no ions are formed. The entire molecule, HCl, is electrically neutral. A partial charge is usually indicated by the Greek letter delta, δ. Thus, a partial positive charge is represented by δ+ and a partial negative charge by δ−. The pair of shared electrons in HCl is closer to the more electronegative chlorine atom than to the hydrogen atom, giving chlorine a partial negative charge with respect to the hydrogen atom. H Cl + δ - δ Hydrogen chloride A scale of relative electronegativities, in which the most electronegative element, fluorine, is assigned a value of 4.0, was developed by the Nobel Laureate (1954 and 1962) Linus Pauling (1901–1994). TABLE 11.5 shows that the relative electronegativity of the nonmetals is high and that of the metals is low.
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