Chemical bond

A chemical bond is a lasting attraction between atoms, ions or molecules that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds. The strength of chemical bonds varies considerably; there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole–dipole interactions, the London dispersion force and hydrogen bonding.

Since opposite charges attract via a simple electromagnetic force, the negatively charged electrons that are orbiting the nucleus and the positively charged protons in the nucleus attract each other. An electron positioned between two nuclei will be attracted to both of them, and the nuclei will be attracted toward electrons in this position. This attraction constitutes the chemical bond. Due to the matter wave nature of electrons and their smaller mass, they must occupy a much larger amount of volume compared with the nuclei, and this volume occupied by the electrons keeps the atomic nuclei in a bond relatively far apart, as compared with the size of the nuclei themselves.

In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. The atoms in molecules, crystals, metals and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure and the bulk properties of matter.

Electron dot
Examples of Lewis dot-style representations of chemical bonds between carbon (C), hydrogen (H), and oxygen (O). Lewis dot diagrams were an early attempt to describe chemical bonding and are still widely used today.

All bonds can be explained by quantum theory, but, in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. More sophisticated theories are valence bond theory which includes orbital hybridization and resonance, and molecular orbital theory which includes linear combination of atomic orbitals and ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances.

Overview of main types of chemical bonds

A chemical bond is an attraction between atoms. This attraction may be seen as the result of different behaviors of the outermost or valence electrons of atoms. These behaviors merge into each other seamlessly in various circumstances, so that there is no clear line to be drawn between them. However it remains useful and customary to differentiate between different types of bond, which result in different properties of condensed matter.

In the simplest view of a covalent bond, one or more electrons (often a pair of electrons) are drawn into the space between the two atomic nuclei. Energy is released by bond formation. This is not as a reduction in potential energy, because the attraction of the two electrons to the two protons is offset by the electron-electron and proton-proton repulsions. Instead, the release of energy (and hence stability of the bond) arises from the reduction in kinetic energy due to the electrons being in a more spatially distributed (i.e. longer de Broglie wavelength) orbital compared with each electron being confined closer to its respective nucleus.[1] These bonds exist between two particular identifiable atoms and have a direction in space, allowing them to be shown as single connecting lines between atoms in drawings, or modeled as sticks between spheres in models.

In a polar covalent bond, one or more electrons are unequally shared between two nuclei. Covalent bonds often result in the formation of small collections of better-connected atoms called molecules, which in solids and liquids are bound to other molecules by forces that are often much weaker than the covalent bonds that hold the molecules internally together. Such weak intermolecular bonds give organic molecular substances, such as waxes and oils, their soft bulk character, and their low melting points (in liquids, molecules must cease most structured or oriented contact with each other). When covalent bonds link long chains of atoms in large molecules, however (as in polymers such as nylon), or when covalent bonds extend in networks through solids that are not composed of discrete molecules (such as diamond or quartz or the silicate minerals in many types of rock) then the structures that result may be both strong and tough, at least in the direction oriented correctly with networks of covalent bonds. Also, the melting points of such covalent polymers and networks increase greatly.

In a simplified view of an ionic bond, the bonding electron is not shared at all, but transferred. In this type of bond, the outer atomic orbital of one atom has a vacancy which allows the addition of one or more electrons. These newly added electrons potentially occupy a lower energy-state (effectively closer to more nuclear charge) than they experience in a different atom. Thus, one nucleus offers a more tightly bound position to an electron than does another nucleus, with the result that one atom may transfer an electron to the other. This transfer causes one atom to assume a net positive charge, and the other to assume a net negative charge. The bond then results from electrostatic attraction between atoms and the atoms become positive or negatively charged ions. Ionic bonds may be seen as extreme examples of polarization in covalent bonds. Often, such bonds have no particular orientation in space, since they result from equal electrostatic attraction of each ion to all ions around them. Ionic bonds are strong (and thus ionic substances require high temperatures to melt) but also brittle, since the forces between ions are short-range and do not easily bridge cracks and fractures. This type of bond gives rise to the physical characteristics of crystals of classic mineral salts, such as table salt.

A less often mentioned type of bonding is metallic bonding. In this type of bonding, each atom in a metal donates one or more electrons to a "sea" of electrons that reside between many metal atoms. In this sea, each electron is free (by virtue of its wave nature) to be associated with a great many atoms at once. The bond results because the metal atoms become somewhat positively charged due to loss of their electrons while the electrons remain attracted to many atoms, without being part of any given atom. Metallic bonding may be seen as an extreme example of delocalization of electrons over a large system of covalent bonds, in which every atom participates. This type of bonding is often very strong (resulting in the tensile strength of metals). However, metallic bonding is more collective in nature than other types, and so they allow metal crystals to more easily deform, because they are composed of atoms attracted to each other, but not in any particularly-oriented ways. This results in the malleability of metals. The cloud of electrons in metallic bonding causes the characteristically good electrical and thermal conductivity of metals, and also their shiny lustre that reflects most frequencies of white light.

History

Early speculations about the nature of the chemical bond, from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. In 1704, Sir Isaac Newton famously outlined his atomic bonding theory, in "Query 31" of his Opticks, whereby atoms attach to each other by some "force". Specifically, after acknowledging the various popular theories in vogue at the time, of how atoms were reasoned to attach to each other, i.e. "hooked atoms", "glued together by rest", or "stuck together by conspiring motions", Newton states that he would rather infer from their cohesion, that "particles attract one another by some force, which in immediate contact is exceedingly strong, at small distances performs the chemical operations, and reaches not far from the particles with any sensible effect."

In 1819, on the heels of the invention of the voltaic pile, Jöns Jakob Berzelius developed a theory of chemical combination stressing the electronegative and electropositive characters of the combining atoms. By the mid 19th century, Edward Frankland, F.A. Kekulé, A.S. Couper, Alexander Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called "combining power", in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert N. Lewis developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond; in Lewis's own words, "An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively."[2]

That same year, Walther Kossel put forward a theory similar to Lewis' only his model assumed complete transfers of electrons between atoms, and was thus a model of ionic bonding. Both Lewis and Kossel structured their bonding models on that of Abegg's rule (1904).

Niels Bohr proposed a model of the atom and a model of the chemical bond. According to his model for a diatomic molecule, the electrons of the atoms of the molecule form a rotating ring whose plane is perpendicular to the axis of the molecule and equidistant from the atomic nuclei. The dynamic equilibrium of the molecular system is achieved through the balance of forces between the forces of attraction of nuclei to the plane of the ring of electrons and the forces of mutual repulsion of the nuclei. The Bohr model of the chemical bond took into account the Coulomb repulsion – the electrons in the ring are at the maximum distance from each other.[3][4]

In 1927, the first mathematically complete quantum description of a simple chemical bond, i.e. that produced by one electron in the hydrogen molecular ion, H2+, was derived by the Danish physicist Øyvind Burrau.[5] This work showed that the quantum approach to chemical bonds could be fundamentally and quantitatively correct, but the mathematical methods used could not be extended to molecules containing more than one electron. A more practical, albeit less quantitative, approach was put forward in the same year by Walter Heitler and Fritz London. The Heitler–London method forms the basis of what is now called valence bond theory. In 1929, the linear combination of atomic orbitals molecular orbital method (LCAO) approximation was introduced by Sir John Lennard-Jones, who also suggested methods to derive electronic structures of molecules of F2 (fluorine) and O2 (oxygen) molecules, from basic quantum principles. This molecular orbital theory represented a covalent bond as an orbital formed by combining the quantum mechanical Schrödinger atomic orbitals which had been hypothesized for electrons in single atoms. The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection (i.e., analytically), but approximations for them still gave many good qualitative predictions and results. Most quantitative calculations in modern quantum chemistry use either valence bond or molecular orbital theory as a starting point, although a third approach, density functional theory, has become increasingly popular in recent years.

In 1933, H. H. James and A. S. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons.[6] With up to 13 adjustable parameters they obtained a result very close to the experimental result for the dissociation energy. Later extensions have used up to 54 parameters and gave excellent agreement with experiments. This calculation convinced the scientific community that quantum theory could give agreement with experiment. However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules.

Bonds in chemical formulas

Because atoms and molecules are three-dimensional, it is difficult to use a single method to indicate orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated in different ways depending on the type of discussion. Sometimes, some details are neglected. For example, in organic chemistry one is sometimes concerned only with the functional group of the molecule. Thus, the molecular formula of ethanol may be written in conformational form, three-dimensional form, full two-dimensional form (indicating every bond with no three-dimensional directions), compressed two-dimensional form (CH3–CH2–OH), by separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons (with the two-dimensional approximate directions) are marked, e.g. for elemental carbon .'C'. Some chemists may also mark the respective orbitals, e.g. the hypothetical ethene−4 anion (\/C=C/\ −4) indicating the possibility of bond formation.

Strong chemical bonds

Typical bond lengths in pm
and bond energies in kJ/mol.

Bond lengths can be converted to Å
by division by 100 (1 Å = 100 pm).
Data taken from University of Waterloo.[7]
Bond Length
(pm)
Energy
(kJ/mol)
H — Hydrogen
H–H 74 436
H–O 96 366
H–F 92 568
H–Cl 127 432
C — Carbon
C–H 109 413
C–C 154 348
C–C= 151
=C–C≡ 147
=C–C= 148
C=C 134 614
C≡C 120 839
C–N 147 308
C–O 143 360
C–F 134 488
C–Cl 177 330
N — Nitrogen
N–H 101 391
N–N 145 170
N≡N 110 945
O — Oxygen
O–O 148 145
O=O 121 498
F, Cl, Br, I — Halogens
F–F 142 158
Cl–Cl 199 243
Br–H 141 366
Br–Br 228 193
I–H 161 298
I–I 267 151

Strong chemical bonds are the intramolecular forces which hold atoms together in molecules. A strong chemical bond is formed from the transfer or sharing of electrons between atomic centers and relies on the electrostatic attraction between the protons in nuclei and the electrons in the orbitals.

The types of strong bond differ due to the difference in electronegativity of the constituent elements. A large difference in electronegativity leads to more polar (ionic) character in the bond.

Ionic bond

Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding, but a difference of electronegativity of over 1.7 is likely to be ionic, and a difference of less than 1.7 is likely to be covalent.[8] Ionic bonding leads to separate positive and negative ions. Ionic charges are commonly between −3e to +3e. Ionic bonding commonly occurs in metal salts such as sodium chloride (table salt). A typical feature of ionic bonds is that the species form into ionic crystals, in which no ion is specifically paired with any single other ion, in a specific directional bond. Rather, each species of ion is surrounded by ions of the opposite charge, and the spacing between it and each of the oppositely charged ions near it, is the same for all surrounding atoms of the same type. It is thus no longer possible to associate an ion with any specific other single ionized atom near it. This is a situation unlike that in covalent crystals, where covalent bonds between specific atoms are still discernible from the shorter distances between them, as measured via such techniques as X-ray diffraction.

Ionic crystals may contain a mixture of covalent and ionic species, as for example salts of complex acids, such as sodium cyanide, NaCN. X-ray diffraction shows that in NaCN, for example, the bonds between sodium cations (Na+) and the cyanide anions (CN) are ionic, with no sodium ion associated with any particular cyanide. However, the bonds between C and N atoms in cyanide are of the covalent type, so that each carbon is strongly bound to just one nitrogen, to which it is physically much closer than it is to other carbons or nitrogens in a sodium cyanide crystal.

When such crystals are melted into liquids, the ionic bonds are broken first because they are non-directional and allow the charged species to move freely. Similarly, when such salts dissolve into water, the ionic bonds are typically broken by the interaction with water, but the covalent bonds continue to hold. For example, in solution, the cyanide ions, still bound together as single CN ions, move independently through the solution, as do sodium ions, as Na+. In water, charged ions move apart because each of them are more strongly attracted to a number of water molecules, than to each other. The attraction between ions and water molecules in such solutions is due to a type of weak dipole-dipole type chemical bond. In melted ionic compounds, the ions continue to be attracted to each other, but not in any ordered or crystalline way.

Covalent bond

Covalent
Nonpolar covalent bonds in methane (CH4). The Lewis structure shows electrons shared between C and H atoms.

Covalent bonding is a common type of bonding, in which two or more atoms share valence electrons more or less equally. The simplest and most common type is a single bond in which two atoms share two electrons. Other types include the double bond, the triple bond, one- and three-electron bonds, the three-center two-electron bond and three-center four-electron bond.

In nonpolar covalent bonds, the electronegativity difference between the bonded atoms is small, typically 0 to 0.3. Bonds within most organic compounds are described as covalent. The figure shows methane (CH4), in which each hydrogen forms a covalent bond with the carbon. See sigma bonds and pi bonds for LCAO-description of such bonding.

Molecules which are formed primarily from non-polar covalent bonds are often immiscible in water or other polar solvents, but much more soluble in non-polar solvents such as hexane.

A polar covalent bond is a covalent bond with a significant ionic character. This means that the two shared electrons are closer to one of the atoms than the other, creating an imbalance of charge. Such bonds occur between two atoms with moderately different electronegativities and give rise to dipole–dipole interactions. The electronegativity difference between the two atoms in these bonds is 0.3 to 1.7.

Single and multiple bonds

A single bond between two atoms corresponds to the sharing of one pair of electrons. The electron density of these two bonding electrons is concentrated in the region between the two atoms, which is the defining quality of a sigma bond.

Pi-Bond
Two p-orbitals forming a pi-bond.

A double bond between two atoms is formed by the sharing of two pairs of electrons, one in a sigma bond and one in a pi bond, with electron density concentrated on two opposite sides of the internuclear axis. A triple bond consists of three shared electron pairs, forming one sigma and two pi bonds.

Quadruple and higher bonds are very rare and occur only between certain transition metal atoms.

Coordinate covalent bond (Dipolar bond)

NH3-BF3-adduct-bond-lengthening-2D
Adduct of ammonia and boron trifluoride

A coordinate covalent bond is a covalent bond in which the two shared bonding electrons are from the same one of the atoms involved in the bond. For example, boron trifluoride (BF3) and ammonia (NH3) form an adduct or coordination complex F3B←NH3 with a B–N bond in which a lone pair of electrons on N is shared with an empty atomic orbital on B. BF3 with an empty orbital is described as an electron pair acceptor or Lewis acid, while NH3 with a lone pair which can be shared is described as an electron-pair donor or Lewis base. The electrons are shared roughly equally between the atoms in contrast to ionic bonding. Such bonding is shown by an arrow pointing to the Lewis acid.

Transition metal complexes are generally bound by coordinate covalent bonds. For example, the ion Ag+ reacts as a Lewis acid with two molecules of the Lewis base NH3 to form the complex ion Ag(NH3)2+, which has two Ag←N coordinate covalent bonds.

Metallic bonding

In metallic bonding, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. The freely-moving or delocalization of bonding electrons leads to classical metallic properties such as luster (surface light reflectivity), electrical and thermal conductivity, ductility, and high tensile strength.

Intermolecular bonding

There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance.

  • A large difference in electronegativity between two bonded atoms will cause a permanent charge separation, or dipole, in a molecule or ion. Two or more molecules or ions with permanent dipoles can interact within dipole-dipole interactions. The bonding electrons in a molecule or ion will, on average, be closer to the more electronegative atom more frequently than the less electronegative one, giving rise to partial charges on each atom, and causing electrostatic forces between molecules or ions.
  • A hydrogen bond is effectively a strong example of an interaction between two permanent dipoles. The large difference in electronegativities between hydrogen and any of fluorine, nitrogen and oxygen, coupled with their lone pairs of electrons cause strong electrostatic forces between molecules. Hydrogen bonds are responsible for the high boiling points of water and ammonia with respect to their heavier analogues.
  • The London dispersion force arises due to instantaneous dipoles in neighbouring atoms. As the negative charge of the electron is not uniform around the whole atom, there is always a charge imbalance. This small charge will induce a corresponding dipole in a nearby molecule; causing an attraction between the two. The electron then moves to another part of the electron cloud and the attraction is broken.
  • A cation–pi interaction occurs between a pi bond and a cation.

Theories of chemical bonding

In the (unrealistic) limit of "pure" ionic bonding, electrons are perfectly localized on one of the two atoms in the bond. Such bonds can be understood by classical physics. The forces between the atoms are characterized by isotropic continuum electrostatic potentials. Their magnitude is in simple proportion to the charge difference.

Covalent bonds are better understood by valence bond (VB) theory or molecular orbital (MO) theory. The properties of the atoms involved can be understood using concepts such as oxidation number, formal charge, and electronegativity. The electron density within a bond is not assigned to individual atoms, but is instead delocalized between atoms. In valence bond theory, bonding is conceptualized as being built up from electron pairs that are localized and shared by two atoms via the overlap of atomic orbitals. The concepts of orbital hybridization and resonance augment this basic notion of the electron pair bond. In molecular orbital theory, bonding is viewed as being delocalized throughout the molecule and apportioned in orbitals that extend throughout the molecule and are adapted to the symmetry properties of the molecule, typically by considering linear combination of atomic orbitals (LCAO). Valence bond theory is more chemically intuitive by being spatially localized, allowing attention to be focused on the parts of the molecule undergoing chemical change. In contrast, molecular orbitals are more "natural" from a quantum mechanical point of view, with orbital energies being physically significant and directly linked to experimental ionization energies from photoelectron spectroscopy. Consequently, valence bond theory and molecular orbital theory are often viewed as competing but complementary frameworks that offer different insights into chemical systems. As approaches for electronic structure theory, both MO and VB methods can both give approximations to any desired level of accuracy, at least in principle. However, at lower levels, the approximations differ, and one approach may be better suited for computations involving a particular system or property than the other.

Unlike the spherically-symmetrical Coulombic forces in pure ionic bonds, covalent bonds are generally directed and anisotropic. These are often classified based on their symmetry with respect to a molecular plane as sigma bonds and pi bonds. In the general case, atoms form bonds that are intermediate between ionic and covalent, depending on the relative electronegativity of the atoms involved. Bonds of this type are known as polar covalent bonds.

References

  1. ^ Rioux, F. (2001). "The Covalent Bond in H2". The Chemical Educator. 6 (5): 288–290. doi:10.1007/s00897010509a.
  2. ^ Lewis, Gilbert N. (1916). "The Atom and the Molecule". Journal of the American Chemical Society. 38 (4): 772. doi:10.1021/ja02261a002. a copy
  3. ^ Бор Н. (1970). Избранные научные труды (статьи 1909–1925). 1. М.: «Наука». p. 133.
  4. ^ Svidzinsky, Anatoly A.; Marlan O. Scully; Dudley R. Herschbach (2005). "Bohr's 1913 molecular model revisited". Proceedings of the National Academy of Sciences. 102 (34[1]): 11985–11988. arXiv:physics/0508161. Bibcode:2005PNAS..10211985S. doi:10.1073/pnas.0505778102. PMC 1186029. PMID 16103360.
  5. ^ Laidler, K. J. (1993). The World of Physical Chemistry. Oxford University Press. p. 346. ISBN 978-0-19-855919-1.
  6. ^ James, H.H.; Coolidge, A S. (1933). "The Ground State of the Hydrogen Molecule". Journal of Chemical Physics. 1 (12): 825–835. Bibcode:1933JChPh...1..825J. doi:10.1063/1.1749252.
  7. ^ "Bond Lengths and Energies". Science.uwaterloo.ca. Retrieved 2013-10-15.
  8. ^ Atkins, Peter; Loretta Jones (1997). Chemistry: Molecules, Matter and Change. New York: W.H. Freeman & Co. pp. 294–295. ISBN 978-0-7167-3107-8.

External links

Bond-dissociation energy

The bond-dissociation energy (BDE, D0, or DH°) is one measure of the strength of a chemical bond A–B. It can be defined as the standard enthalpy change when A–B is cleaved by homolysis to give fragments A and B, which are usually radical species. The enthalpy change is temperature dependent, and the bond-dissociation energy is often defined to be the enthalpy change of the homolysis at 0 K (absolute zero), although the enthalpy change at 298 K (standard conditions) is also a frequently encountered parameter. As a typical example, the bond-dissociation energy for one of the C–H bonds in ethane (C2H6) is defined as the standard enthalpy change of the process

CH3CH2–H → CH3CH2• + H•,

DH°298 (CH3CH2–H) = ΔH° = 101.1(4) kcal/mol = 423.0 ± 1.7 kJ/mol = 4.40(2) eV (per bond).To convert a molar BDE to the energy needed to dissociate the bond per molecule, the conversion factor 23.060 kcal/mol (96.485 kJ/mol) for each eV can be used.

A variety of experimental techniques, including spectrometric determination of energy levels, generation of radicals by pyrolysis or photolysis, measurements of chemical kinetics and equilibrium, and various calorimetric and electrochemical methods have been used to measure bond dissociation energy values. Nevertheless, bond dissociation energy measurements are challenging and are subject to considerable error. The majority of currently known values are accurate to within ± 1 or 2 kcal/mol. Moreover, values measured in the past, especially before the 1970s, can be especially unreliable and have been subject to revisions on the order of 10 kcal/mol (e.g., benzene C-H bonds, from 103 kcal/mol in 1965 to the modern accepted value of 112.9(5) kcal/mol). Even in modern times (between 1990 and 2004), the O-H bond of phenol has been reported to be anywhere from 85.8 to 91.0 kcal/mol. On the other hand, the bond dissociation energy of H2 at 298 K has been measured to high precision and accuracy: DH°298 (H–H) = 104.1539(1) kcal/mol.

Bond dipole moment

The bond dipole moment uses the idea of electric dipole moment to measure the polarity of a chemical bond within a molecule. It occurs whenever there is a separation of positive and negative charges. The bond dipole μ is given by:

.

The bond dipole is modeled as +δ — δ- with a distance d between the partial charges +δ and δ-. It is a vector, parallel to the bond axis, pointing from minus to plus, as is conventional for electric dipole moment vectors.

Chemists often draw the vector pointing from plus to minus. This vector can be physically interpreted as the movement undergone by electrons when the two atoms are placed a distance d apart and allowed to interact, the electrons will move from their free state positions to be localised more around the more electronegative atom.

The SI unit for electric dipole moment is the coulomb-meter. This is too large to be practical on the molecular scale. Bond dipole moments are commonly measured in debyes, represented by the symbol D, which is obtained by measuring the charge in units of 10−10 statcoulomb and the distance d in Angstroms. Note that 10−10 statcoulomb is 0.208 units of elementary charge, so 1.0 debye results from an electron and a proton separated by 0.208 Angstrom. A useful conversion factor is 1 D = 3.335 64×1030C m.

For diatomic molecules there is only one (single or multiple) bond so the bond dipole moment is the molecular dipole moment, with typical values in the range of 0 to 11 D. At one extreme, a symmetrical molecule such as chlorine, Cl
2
, has zero dipole moment, while near the other extreme, gas phase potassium bromide, KBr, which is highly ionic, has a dipole moment of 10.5 D.

For polyatomic molecules there is more than one bond, and the total molecular dipole moment may be approximated as the vector sum of individual bond dipole moments. Often bond dipoles are obtained by the reverse process: a known total dipole of a molecule can be decomposed into bond dipoles. This is done to transfer bond dipole moments to molecules that have the same bonds, but for which the total dipole moment is not yet known. The vector sum of the transferred bond dipoles gives an estimate for the total (unknown) dipole of the molecule.

Bond energy

In chemistry, bond energy (E) or bond enthalpy (H) is the measure of bond strength in a chemical bond. IUPAC defines bond energy as the average value of the gas-phase bond dissociation energies (usually at a temperature of 298 K) for all bonds of the same type within the same chemical species. For example, the carbon–hydrogen bond energy in methane H(C–H) is the enthalpy change involved with breaking up one molecule of methane into a carbon atom and four hydrogen radicals, divided by 4. Tabulated bond energies are generally values of bond energies averaged over a number of selected typical chemical species containing that type of bond. Bond energy (E) or bond enthalpy (H) should not be confused with bond-dissociation energy. Bond energy is the average of all the bond-dissociation energies in a molecule, and will show a different value for a given bond than the bond-dissociation energy would. This is because the energy required to break a single bond in a specific molecule differs for each bond in that molecule. For example, methane has four C–H bonds and the bond-dissociation energies are 435 kJ/mol for D(CH3–H), 444 kJ/mol for D(CH2–H), 444 kJ/mol for D(CH–H) and 339 kJ/mol for D(C–H). Their average, and hence the bond energy, is 414 kJ/mol, even though not a single bond required specifically 414 kJ/mol to be broken.

Bond length

In molecular geometry, bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types, relatively independent of the rest of the molecule.

Bond order

Bond order is the number of chemical bonds between a pair of atoms. For example, in diatomic nitrogen N≡N the bond order is 3, in acetylene H−C≡C−H the bond order between the two carbon atoms is also 3, and the C−H bond order is 1. Bond order gives an indication of the stability of a bond. Isoelectronic species have same bond number.In molecules which have resonance or nonclassical bonding, bond order may not be an integer. In benzene, the delocalized molecular orbitals contain 6 pi electrons over six carbons essentially yielding half a pi bond together with the sigma bond for each pair of carbon atoms, giving a calculated bond order of 1.5. Furthermore, bond orders of 1.1, for example, can arise under complex scenarios and essentially refer to bond strength relative to bonds with order 1.

Carbon–fluorine bond

The carbon–fluorine bond is a polar covalent bond between carbon and fluorine that is a component of all organofluorine compounds. It is the fourth strongest single bond in organic chemistry—behind the B-F single bond, Si-F single bond and the H-F single bond, and relatively short—due to its partial ionic character. The bond also strengthens and shortens as more fluorines are added to the same carbon on a chemical compound. As such, fluoroalkanes like tetrafluoromethane (carbon tetrafluoride) are some of the most unreactive organic compounds.

Covalent bond

A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding. For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration.

Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, agostic interactions, bent bonds, and three-center two-electron bonds. The term covalent bond dates from 1939. The prefix co- means jointly, associated in action, partnered to a lesser degree, etc.; thus a "co-valent bond", in essence, means that the atoms share "valence", such as is discussed in valence bond theory.

In the molecule H2, the hydrogen atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities. Thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons over more than two atoms is said to be delocalized.

Homolysis (chemistry)

In chemistry, homolysis (from Greek ὅμοιος, homoios, "equal," and λύσις, lusis, "loosening") or homolytic fission is chemical bond dissociation of a molecule by a process where each of the fragments retains one of the originally bonded electrons. During homolytic fission of a neutral molecule with an even number of electrons, two free radicals will be generated. That is, the two electrons involved in the original bond are distributed between the two fragment species. The energy involved in this process is called bond dissociation energy. Bond cleavage is also possible by a process called heterolysis.

Because the relatively high energy required to break bonds in this manner, homolysis only occurs under certain circumstances:

Ultraviolet radiation (e.g. sunlight causing sunburns)

Heat

Certain intramolecular bonds, such as the O–O bond of a peroxide, are weak enough to spontaneously homolytically dissociate with a small amount of heat.

High temperatures in the absence of oxygen (pyrolysis) can induce homolytic elimination of carbon compounds.

Hydrolase

Hydrolase is a class of enzyme that is commonly used as biochemical catalysts that use water to break a chemical bond. This results in a division of a larger molecule to smaller molecules. Some common examples of hydrolase enzymes are esterases including lipases, phosphatases, glycosidases, peptidases, and nucleosidases. Esterases cleave ester bonds in lipids and phosphatases cleave phosphate groups off molecules. An example of crucial esterase is the acetylcholine esterase, which assists in transforming the neuron impulse into acetic acid after it the hydrolase breaks the acetylcholine into choline and acetic acid. Acetic acid is an important metabolite in the body, which becomes a nice intermediate for other reactions such as glycolysis. Lipases hydrolyze glycerides. Glycosidases cleave sugar molecules off carbohydrates and peptidases hydrolyze peptide bonds. Nucleosidases hydrolyze the bonds of nucleotides.Hydrolase enzymes are important for the body because they have degradative properties. In lipids, lipases contribute to the breakdown of fats and lipoproteins and other larger molecules into smaller molecules like fatty acids and glycerol. Fatty acids and other small molecules are used for synthesis and as a source of energy.In biochemistry, a hydrolase is an enzyme that catalyzes the hydrolysis of a chemical bond. For example, an enzyme that catalyzes the following reaction is a hydrolase:

A–B + H2O → A–OH + B–H

Ligase

In biochemistry, a ligase is an enzyme that can catalyze the joining of two large molecules by forming a new chemical bond, usually with accompanying hydrolysis of a small pendant chemical group on one of the larger molecules or the enzyme catalyzing the linking together of two compounds, e.g., enzymes that catalyze joining of C-O, C-S, C-N, etc. In general, a ligase catalyzes the following reaction:

Ab + C → A–C + bor sometimes

Ab + cD → A–D + b + c + d + e + fwhere the lowercase letters can signify the small, dependent groups. Ligase can join two complementary fragments of nucleic acid and repair single stranded breaks that arise in double stranded DNA during replication.

Linus Pauling

Linus Carl Pauling (; February 28, 1901 – August 19, 1994) was an American chemist, biochemist, peace activist, author, educator, and husband of American human rights activist Ava Helen Pauling. He published more than 1,200 papers and books, of which about 850 dealt with scientific topics. New Scientist called him one of the 20 greatest scientists of all time, and as of 2000, he was rated the 16th most important scientist in history.Pauling was one of the founders of the fields of quantum chemistry and molecular biology. His contributions to the theory of the chemical bond include the concept of orbital hybridisation and the first accurate scale of electronegativities of the elements. Pauling also worked on the structures of biological molecules, and showed the importance of the alpha helix and beta sheet in protein secondary structure. Pauling's approach combined methods and results from X-ray crystallography, molecular model building and quantum chemistry. His discoveries inspired the work of James Watson, Francis Crick, and Rosalind Franklin on the structure of DNA, which in turn made it possible for geneticists to crack the DNA code of all organisms.In his later years he promoted nuclear disarmament, as well as orthomolecular medicine, megavitamin therapy, and dietary supplements. None of the latter have gained much acceptance in the mainstream scientific community.For his scientific work, Pauling was awarded the Nobel Prize in Chemistry in 1954. For his peace activism, he was awarded the Nobel Peace Prize in 1962. He is one of four individuals to have won more than one Nobel Prize (the others being Marie Curie, John Bardeen and Frederick Sanger). Of these, he is the only person to have been awarded two unshared Nobel Prizes, and one of two people to be awarded Nobel Prizes in different fields, the other being Marie Curie.

Nucleophile

Nucleophile is a chemical species that donates an electron pair to an electrophile to form a chemical bond in relation to a reaction. All molecules or ions with a free pair of electrons or at least one pi bond can act as nucleophiles. Because nucleophiles donate electrons, they are by definition Lewis bases.

Nucleophilic describes the affinity of a nucleophile to the nuclei. Nucleophilicity, sometimes referred to as nucleophile strength, refers to a substance's nucleophilic character and is often used to compare the affinity of atoms. Neutral nucleophilic reactions with solvents such as alcohols and water are named solvolysis. Nucleophiles may take part in nucleophilic substitution, whereby a nucleophile becomes attracted to a full or partial positive charge.

Organoactinide chemistry

Organoactinide chemistry is the science exploring the properties, structure and reactivity of organoactinide compounds, which are organometallic compounds containing a carbon to actinide chemical bond.

Organoscandium chemistry

Organoscandium chemistry is the chemistry of organometallic compounds containing a carbon to scandium chemical bond. The interest in organoscandium compounds is mostly academic but several compound classes find practical application in catalysis, especially in polymerization. A common precursor is scandium chloride.

As with the other elements in group 3 – e.g. yttrium, forming organoyttrium compounds – and the lanthanides, the dominant oxidation state for scandium in organometallic compounds is +3 (electron configuration [Ar] 3d14s2). The members of this group also have large ionic radii with vacant s,p and d orbitals (88 pm for Sc3+ compared to 67 pm for Al3+) and as a result they behave as hard Lewis acids and tend to have high coordination numbers of 9 to 12. The metal to ligand chemical bond is largely ionic.

Potassium bifluoride

Potassium bifluoride is the inorganic compound with the formula KHF2. This colourless salt consists of the potassium cation and the bifluoride (HF2−) anion. The salt is used in etchant for glass. Sodium bifluoride is related and is also of commercial use as an etchant as well as in cleaning products.

Staggered conformation

In organic chemistry, a staggered conformation is a chemical conformation of an ethane-like moiety abcX–Ydef in which the substituents a, b, and c are at the maximum distance from d, e, and f. This requires the torsion angles to be 60°.Eliel, Ernest L.; Wilen, Samuel H. (1994). Stereochemistry of Organic Compounds. Wiley. p. 1207. ISBN 978-0-471-01670-0.Such a conformation exists in any open chain single chemical bond connecting two sp3-hybridised atoms, and is normally a conformational energy minimum. For some molecules such as those of n-butane, there can be special versions of staggered conformations called gauche and anti; see first Newman projection diagram in Conformational isomerism.

Crystal structures

The staggered/eclipsed configurations distinguish different crystalline structures of eg. cubic/hexagonal boron nitride, and diamond/lonsdaleite.

Three-center two-electron bond

A 3-center 2-electron (3c–2e) bond is an electron-deficient chemical bond where three atoms share two electrons. The combination of three atomic orbitals form three molecular orbitals: one bonding, one non-bonding, and one anti-bonding. The two electrons go into the bonding orbital, resulting in a net bonding effect and constituting a chemical bond among all three atoms. In many common bonds of this type, the bonding orbital is shifted towards two of the three atoms instead of being spread equally among all three. An example of a 3c-2e bond is the trihydrogen cation H+3. This type of bond is also called banana bond.

Triple bond

A triple bond in chemistry is a chemical bond between two atoms involving six bonding electrons instead of the usual two in a covalent single bond. The most common triple bond, that between two carbon atoms, can be found in alkynes.

Other functional groups containing a triple bond are cyanides and isocyanides. Some diatomic molecules, such as dinitrogen and carbon monoxide, are also triple bonded. In skeletal formula the triple bond is drawn as three parallel lines (≡) between the two connected atoms.Triple bonds are stronger than the equivalent single bonds or double bonds, with a bond order of three.

Valence bond theory

In chemistry, valence bond (VB) theory is one of two basic theories, along with molecular orbital (MO) theory, that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the atomic orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule.

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