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Encyclopedia > Chemical bond

A chemical bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. In practice, however, chemists usually rely on quantum theory or qualitative descriptions that are less rigorous but more easily explained to describe chemical bonding. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. Molecules, crystals, and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure of matter. Properties For other meanings of Atom, see Atom (disambiguation). ... 3D (left and center) and 2D (right) representations of the terpenoid molecule atisane. ... Look up chemical compound in Wiktionary, the free dictionary. ... Quantum electrodynamics (QED) is a relativistic quantum field theory of electrodynamics. ... For a less technical and generally accessible introduction to the topic, see Introduction to quantum mechanics. ... 3D (left and center) and 2D (right) representations of the terpenoid molecule atisane. ... For other uses, see Crystal (disambiguation). ... This article does not cite any references or sources. ...


Bonds vary widely in their strength. Generally covalent and ionic bonds are often described as "strong", whereas hydrogen bonds and van der Waals' bonds are generally considered to be "weak". Care should be taken because the strongest of the "weak" bonds can be stronger than the weakest of the "strong" bonds. Covalent redirects here. ... Sodium and chlorine bonding ionically to form sodium chloride. ... An example of a quadruple hydrogen bond between a self-assembled dimer complex reported by Meijer and coworkers. ... In chemistry, the term van der Waals forces (sometimes called London dispersion forces) refers to a particular class of intermolecular forces. ...

Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.

Contents

Image File history File links Electron_dot. ... Image File history File links Electron_dot. ... Lewis structures, also called electron-dot structures or electron-dot diagrams, are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. ... For other uses, see Carbon (disambiguation). ... This article is about the chemistry of hydrogen. ... General Name, symbol, number oxygen, O, 8 Chemical series nonmetals, chalcogens Group, period, block 16, 2, p Appearance colorless (gas) pale blue (liquid) Standard atomic weight 15. ...

Overview

Remembering that opposite charges attract, and that the electrons orbiting the nucleus are negatively charged and protons in the nucleus are positively charged, then imagine two atoms near each other which form a covalent bond. Properties The electron (also called negatron, commonly represented as e−) is a subatomic particle. ... For alternative meanings see proton (disambiguation). ... Covalent redirects here. ...


In the simplest view of a so-called covalent bond, one or more electrons—often a pair as in this example—is drawn into the space between the two atomic nuclei. Here the negatively charged electrons are attracted to the positive charges of both nuclei, instead of just their own. This overcomes the repulsion between the two positively charged nuclei of the two atoms and so this overwhelming attraction holds the two nuclei in a relatively fixed configuration of equilibrium, even though they will still vibrate at equilibrium position. In summary, covalent bonding involves sharing of electrons in which the positively charged nuclei of two or more atoms simultaneously attract the negatively charged electrons that are being shared.


In a simplified view of an ionic bond, the positive charge of one of the nuclei overwhelms the positive charge of the other nucleus, thus effectively transferring an electron from one atom to another, causing 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. This article is about the electrically charged particle. ...


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 the linear combination of atomic orbitals molecular orbital method which includes ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances. The bonding in carbon dioxide The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, similar to the electronic configuration of a noble gas. ... Valence shell electron pair repulsion theory (VSEPR) (1957) is a model in chemistry that aims to generally represent the shapes of individual molecules [1] . To achieve this, it is necessary to construct a valid Lewis structure that shows all of the bonds within the molecule and the locations of lone... In chemistry, valence bond theory explains the nature of a chemical bond in a molecule in terms of atomic valencies. ... In chemistry, hybridisation is the mixing of atomic orbitals to form new orbitals suitable for bonding. ... For other uses, see Resonance (disambiguation). ... This article may be too technical for most readers to understand. ... Ligand field theory was developed during the thirties and fourties of the twentieth century as an expansion of the electrostatic crystal field theory, which offered a good description of the electronic structure of metal ions in coordination complexes but was not able to provide a proper explanation for their bonding. ... Electrostatics (also known as static electricity) is the branch of physics that deals with the phenomena arising from what seem to be stationary electric charges. ...


History

Early speculations into 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, Issac 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: Portrait of Monsieur Lavoisier and his Wife, by Jacques-Louis David The history of chemistry is long and convoluted. ... In chemistry, the history of the molecule traces the origins of the concept or idea of the existence, in nature, of a bonded structure of two or more atoms, according to which the structures of the universe are built. ... Chemical species is a common, general name for atoms, molecules, molecular fragments and ions as entities being subjected to a chemical process or to a measurement. ... Chemical affinity results from electronic properties by which dissimilar substances are capable of forming chemical compounds. ... Sir Isaac Newton in Knellers portrait of 1689. ... Opticks or a treatise of the reflections, refractions, inflections and colours of light Opticks is a book written by English physicist Isaac Newton that was released to the public in 1704. ... Properties For other meanings of Atom, see Atom (disambiguation). ... For other uses, see Force (disambiguation). ...

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 character of the combining atoms. By the mid 19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. 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: For other uses, see Force (disambiguation). ... A copper-zinc Voltaic pile A Voltaic pile on display in the Tempio Voltiano The Voltaic pile is the first modern electric battery, invented by Alessandro Volta in 1800. ... Friherre Jöns Jakob Berzelius (August 20, 1779 – August 7, 1848) was a Swedish chemist. ... Sir Edward Frankland (January 18, 1825 – August 9, 1899) was an English chemist. ... Adolph Wilhelm Hermann Kolbe (September 27, 1818 – November 25, 1884) was a chemist. ... In chemistry, radicals (often referred to as free radicals) are atomic or molecular species with unpaired electrons on an otherwise open shell configuration. ... In chemistry, valence, also known as valency or valency number, is a measure of the number of chemical bonds formed by the atoms of a given element. ... Lewis in the Berkeley Lab Gilbert Newton Lewis (October 23, 1875-March 23, 1946) was a famous American physical chemist. ... Covalent redirects here. ... A chemical bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. ... Covalently bonded hydrogen and carbon in a molecule of methane. ... Covalent bonding is a form of chemical bonding characterized by the sharing of one or more pairs of electrons between atoms, in order to produce a mutual attraction, which holds the resultant molecule together. ... Covalent bonding is a form of chemical bonding characterized by the sharing of one or more pairs of electrons between atoms, in order to produce a mutual attraction, which holds the resultant molecule together. ...

In Lewis' own words: Image File history File links Lewis-bond. ...

An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively.

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 polar bonds. Both Lewis and Kossel structured their bonding models on that of Abegg's rule (1904). Walther Ludwig Julius Kossel (January 4, 1888 in Berlin, Germany – 22 May 1956 in Tübingen, Germany) was a German physicist known for his theory of the chemical bond (ionic bond/octet rule), Sommerfeld-Kossel displacement law of atomic spectra, the Kossel-Stranski model for crystal growth, and the Kossel... Chemical polarity refers how polar a chemical bond is. ... In chemistry, Abegg’s rule states that the difference between the maximum positive and negative valence of an element is frequently eight. ...


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 Oyvind Burrau.[1] 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 a orbitals 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 preditions 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. Walter Heinrich Heitler (02. ... Fritz Wolfgang London (March 7, 1900–March 30, 1954) was a German-born American physicist for whom the London force is named. ... In chemistry, valence bond theory explains the nature of a chemical bond in a molecule in terms of atomic valencies. ... This article may be too technical for most readers to understand. ... John Edward Lennard-Jones (October 27, 1894 - November 1, 1954) was a mathematician who held a chair of theoretical physics at Bristol University, and then a chair of theoretical science at Cambridge University. ... Distinguished from fluorene and fluorone. ... General Name, symbol, number oxygen, O, 8 Chemical series nonmetals, chalcogens Group, period, block 16, 2, p Appearance colorless (gas) pale blue (liquid) Standard atomic weight 15. ... In chemistry, a molecular orbital is a region in which an electron may be found in a molecule. ... For a non-technical introduction to the topic, please see Introduction to quantum mechanics. ... Quantum chemistry is a branch of theoretical chemistry, which applies quantum mechanics and quantum field theory to address issues and problems in chemistry. ... Density functional theory (DFT) is a quantum mechanical method used in physics and chemistry to investigate the electronic structure of many-body systems, in particular molecules and the condensed phases. ...


In 1935, 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.[2] 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 give excellent agreement with experiment. 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.


Valence bond theory

Main article: Valence bond theory

In the year 1927, valence bond theory was formulated which argued essentially that a chemical bond forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of system energy lowering effects. In 1931, building on this theory, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were already generally known: In chemistry, valence bond theory explains the nature of a chemical bond in a molecule in terms of atomic valencies. ... Valence is a scientific term in chemistry to describe electrons in the outermost orbital. ... In chemistry, an atomic orbital is the region in which an electron may be found around a single atom. ... Linus Carl Pauling (February 28, 1901 – August 19, 1994) was an American quantum chemist and biochemist. ...

1. The electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
2. The spins of the electrons have to be opposed.
3. Once paired, the two electrons cannot take part in additional bonds.

His last three rules were new:

4. The electron-exchange terms for the bond involves only one wave function from each atom.
5. The available electrons in the lowest energy level form the strongest bonds.
6. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.

Building on this article, Pauling’s 1939 textbook: On the Nature of the Chemical Bond would become what some have called the “bible” of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to address adequately the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960's and 1970's as molecular orbital theory grew in popularity and was implemented in many large computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory into computer programs have been largely solved and valence bond theory has seen a resurgence.


Molecular orbital theory

Molecular orbital theory (MO) uses a linear combination of atomic orbitals to form molecular orbitals which cover the whole molecule. These are often divided into bonding orbitals, anti-bonding orbitals, and non-bonding orbitals. A molecular orbital is merely a Schrödinger orbital which includes several, but often only two nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being between nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an anti-bonding orbital and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly atomic orbitals) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength. In chemistry, molecular orbital theory (MO theory) is a method for determining molecular structure in which electrons are not assigned to individual bonds between atoms, but are treated as moving under the influence of the nuclei in the whole molecule. ... In chemistry, a molecular orbital is a region in which an electron may be found in a molecule. ... In chemistry, an atomic orbital is the region in which an electron may be found around a single atom. ... Antibonding (or anti-bonding) is a type of chemical bonding. ... In chemistry, a molecular orbital is a region in which an electron may be found in a molecule. ... Antibonding (or anti-bonding) is a type of chemical bonding. ... An atomic orbital is the description of the behavior of an electron in an atom according to quantum mechanics. ...


Comparison of valence bond and molecular orbital theory

In some respects valence bond theory is superior to molecular orbital theory. When applied to the simplest two-electron molecule, H2, valence bond theory, even at the simplest Heitler-London approach, gives a much closer approximation to the bond energy, and it provides a much more accurate representation of the behavior of the electrons as chemical bonds are formed and broken. In contrast simple molecular orbital theory predicts that the hydrogen molecule dissociates into a linear superposition of hydrogen atoms and positive and negative hydrogen ions, a completely unphysical result. This explains in part why the curve of total energy against interatomic distance for the valence bond method lies above the curve for the molecular orbital method at all distances and most particularly so for large distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for F2, where the minimum energy of the curve with molecular orbital theory is still higher in energy than the energy of two F atoms. In chemistry, bond energy (E) is a measure of bond strength in a chemical bond. ...


The concepts of hybridization are so versatile, and the variability in bonding in most organic compounds is so modest, that valence bond theory remains an integral part of the vocabulary of organic chemistry. However, the work of Friedrich Hund, Robert Mulliken, and Gerhard Herzberg showed that molecular orbital theory provided a more appropriate description of the spectroscopic, ionization and magnetic properties of molecules. The deficiencies of valence bond theory became apparent when hypervalent molecules (e.g. PF5) were explained without the use of d orbitals that were crucial to the bonding hybridisation scheme proposed for such molecules by Pauling. Metal complexes and electron deficient compounds (e.g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made. Carl von Weizacker & Friedrich Hund, Goettingen DPI Friedrich Hund (February 4, 1896 - March 31, 1997) : German physicist known for his work on atoms and molecules. ... Robert Sanderson Mulliken (June 7, 1896-October 31, 1986) was an American physicist and chemist, primarily responsible for the elaboration of the molecular orbital method of computing the structure of molecules. ... Gerhard Herzberg (December 25, 1904 – March 3, 1999) was a pioneering theoretical chemist. ... Synthesis of copper(II)-tetraphenylporphine, a metal complex, from tetraphenylporphine and copper(II) acetate monohydrate. ... A compound that is electron deficient has too few valence electrons for the connections between atoms to be described as covalent bonds. ... Diborane is a colorless gas at room temperature with a repulsive, sweet odor. ...


In the 1930s the two methods strongly competed until it was realised that they are both approximations to a better theory. If we take the simple valence bond structure and mix in all possible covalent and ionic structures arising from a particular set of atomic orbitals, we reach what is called the full configuration interaction wave function. If we take the simple molecular orbital description of the ground state and combine that function with the functions describing all possible excited states using unoccupied orbitals arising from the same set of atomic orbitals, we also reach the full configuration interaction wavefunction. It can be then seen that the simple molecular orbital approach gives too much weight to the ionic structures, while the simple valence bond approach gives too little. This can also be described as saying that the molecular orbital approach is too delocalised, while the valence bond approach is too localised.


The two approaches are now regarded as complementary, each providing its own insights into the problem of chemical bonding. Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. However better valence bond programs are now available. Quantum chemistry is a branch of theoretical chemistry, which applies quantum mechanics and quantum field theory to address issues and problems in chemistry. ...


Bonds in chemical formulas

The 3-dimensionality of atoms and molecules makes it difficult to use a single technique for indicating orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH3–CH2–OH), 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 2-dimensionalized approximate directions) are marked, i.e. for elemental carbon .'C'. Some chemists may also mark the respective orbitals, i.e. the hypothetical ethene−4 anion (/C=C/ −4) indicating the possibility of bond formation. A chemical formula (also called molecular formula) is a concise way of expressing information about the atoms that constitute a particular chemical compound. ... Organic chemistry is a specific discipline within chemistry which involves the scientific study of the structure, properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds consisting primarily of carbon and hydrogen, which may contain any number of other elements, including nitrogen, oxygen, halogens as well... In organic chemistry, functional groups (or moieties) are specific groups of atoms within molecules, that are responsible for the characteristic chemical reactions of those molecules. ... Alcoholic beverages An alcoholic beverage is a drink containing ethanol, commonly known as alcohol, although in chemistry the definition of alcohol includes many other compounds. ... Conformational isomerism is the phenomenon of molecules with the same structural formula but different conformations (conformers) of atoms about a rotating bond. ...


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 [1].
Bond Length
(pm)
Energy
(kJ/mol)
H — Hydrogen
H–H 74 436
H–C 109 413
H–N 101 391
H–O 96 366
H–F 92 568
H–Cl 127 432
H–Br 141 366
C — Carbon
C–H 109 413
C–C 154 348
C=C 134 614
C≡C 120 839
C–N 147 308
C–O 143 360
C–F 135 488
C–Cl 177 330
C–Br 194 288
C–I 214 216
C–S 182 272
N — Nitrogen
N–H 101 391
N–C 147 308
N–N 145 170
N≡N 110 945
O — Oxygen
O–H 96 366
O–C 143 360
O–O 148 145
O=O 121 498
F, Cl, Br, I — Halogens
F–H 92 568
F–F 142 158
F–C 135 488
Cl–H 127 432
Cl–C 177 330
Cl–Cl 199 243
Br–H 141 366
Br–C 194 288
Br–Br 228 193
I–H 161 298
I–C 214 216
I–I 267 151
S — Sulfur
C–S 182 272

These chemical bonds are intramolecular forces, which hold atoms together in molecules. In the simplistic localized view of bonding, the number of electrons participating in a bond (or located in a bonding orbital) is typically multiples of two, four, or six, respectively. Even numbers are common because electrons enjoy lower energy states, if paired. Substantially more advanced bonding theories have shown that bond strength is not always a whole number, depending on the distribution of electrons to each atom involved in a bond. For example, the carbons in benzene are connected to each other with about 1.5 bonds, and the two atoms in nitric oxide NO, are connected with about 2.5 bonds. Quadruple bonds are also well known. The type of strong bond depends on the difference in electronegativity and the distribution of the electron orbital paths available to the atoms that are bonded. The larger the difference in electronegativity, the more an electron is attracted to a particular atom involved in the bond, and the more "ionic" properties the bond is said to have ("ionic" means the bond electron(s) are unequally shared). The smaller the difference in electronegativity, the more covalent properties (full sharing) the bond has. In molecular geometry, bond length or bond distance is the distance between two bonded atoms in a molecule. ... An Ã¥ngström or aangstroem (the official transliteration), or angstrom (symbol Ã…) is a non-SI unit of length that is internationally recognized, equal to 0. ... This article is about the chemistry of hydrogen. ... For other uses, see Carbon (disambiguation). ... General Name, symbol, number nitrogen, N, 7 Chemical series nonmetals Group, period, block 15, 2, p Appearance colorless gas Standard atomic weight 14. ... General Name, symbol, number oxygen, O, 8 Chemical series nonmetals, chalcogens Group, period, block 16, 2, p Appearance colorless (gas) pale blue (liquid) Standard atomic weight 15. ... This article is about the chemical series. ... This article is about the chemical element. ... In science, a molecule is the smallest particle of a pure chemical substance that still retains its chemical composition and properties. ... In chemistry, bond strength is measured between two atoms joined in a chemical bond. ... For benzine, see petroleum ether. ... R-phrases , , , , S-phrases , , , Except where noted otherwise, data are given for materials in their standard state (at 25 Â°C, 100 kPa) Infobox disclaimer and references Nitric oxide or Nitrogen monoxide is a chemical compound with chemical formula NO. This gas is an important signaling molecule in the body of... Covalently bonded hydrogen and carbon in a molecule of methane. ... Electronegativity is a measure of the ability of an atom or molecule to attract electrons in the context of a chemical bond. ...


Covalent bond

Main article: Covalent bond

Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such bonding. Covalent redirects here. ... Benzene is the simplest of the arenes, a family of organic compounds An organic compound is any member of a large class of chemical compounds whose molecules contain carbon. ... Electron atomic and molecular orbitals, showing among others the sigma bond of two s-orbitals and a sigma bond of two p-orbitals In chemistry, sigma bonds (σ bonds) are a type of covalent chemical bond. ... Electron atomic and molecular orbitals, showing a Pi-bond at the bottom right of the picture. ...


Polar covalent bond

Main article: Polar covalent bond

Polar covalent bonding is intermediate in character between a covalent and an ionic bond. A polar covalent bond is a form of covalent bonding that happens when atoms of two different elements with different electronegativities bond resulting in an unequal sharing of electrons. ...


Ionic bond

Main article: Ionic bond

Ionic bonding is a type of electrostatic interaction between atoms which have an electronegativity difference of over 1.6 (this limit is a convention).[citation needed] These form in a solution between two ions after the excess of the solvent is removed. Ionic charges are commonly between −3e to +7e Sodium and chlorine bonding ionically to form sodium chloride. ... ... The elementary charge (symbol e or sometimes q) is the electric charge carried by a single proton, or equivalently, the negative of the electric charge carried by a single electron. ... The elementary charge (symbol e or sometimes q) is the electric charge carried by a single proton, or equivalently, the negative of the electric charge carried by a single electron. ...


Coordinate covalent bond

Coordinate covalent bonding, sometimes referred to as dative bonding, is a kind of covalent bonding, in which the covalent bonding electrons originate solely from one of the atoms, the electron-pair donor or Lewis base but are approximately equally shared in the formation of a covalent bond. This concept is somewhat fading as chemists increasingly embrace molecular orbital theory. Examples of coordinate covalent bonding occur in nitrones and ammonia borane. The arrangement is different from an ionic bond in that the electronegativity difference is small, resulting in covalency. They are shown by an arrow unlike other bonds . This arrow shows its head towards the electron acceptor or lewis acid and its tail towards the lewis base. This bond is seen in ammonium. A coordinate covalent bond (also known as dative bond) is a description of covalent bonding in many kinds of compounds. ... A nitrone is the N-oxide of an imine and a functional group in organic chemistry. ... Ammonia borane, NH3BH3, is an inorganic hydride of nitrogen and boron. ...


Bent bonds

Main article: Bent bond

Bent bonds, also known as banana bonds, are bonds in strained or otherwise sterically hindered molecules those binding orbitals are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds. This article is about chemistry. ... This article is about chemistry. ... Steric effects are the interaction of molecules dictated by their shape and/or spatial relationships. ...


3c-2e and 3c-4e bonds

In three-center two-electron bonds three atoms share two electrons in bonding. This type of bonding occurs in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape (shown as a more sharply angled section in the stick model at right), with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms. Three-center four-electron bonds also exist which explain the bonding in hypervalent molecules. In certain cluster compounds so-called four-center two-electron bonds also have been postulated. A three-center two-electron bond is an electron deficient chemical bond where three atoms share two electrons. ... Diborane is a colorless gas at room temperature with a repulsive, sweet odor. ... The 3-center-4-electron bond is a model used to explain bonding in hypervalent molecules such as phosphorus pentafluoride, sulfur hexafluoride, the xenon fluorides, and the hydrogen difluoride ion. ... A hypervalent molecule is a molecule that contains one or more typical elements (group 1, 2, 13-18) formally bearing more than eight electrons in their valence shells. ... A four-center two-electron bond is a type of chemical bond in which four atoms share two electrons in bonding which is unusual because in ordinary chemical bonds two atoms share two electrons (2c-2e bonding). ...


One- and three-electron bonds

Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the hydrogen molecular cation, H2+. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li2+ than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects. [3] In chemistry, radicals (often referred to as free radicals) are atomic or molecular species with unpaired electrons on an otherwise open shell configuration. ... Dilithium is a fictional crystalline mineral in the universe of Star Trek. ...


The simplest example of three-electron bonding can be found in the helium dimer cation, He2+, and can also be considered a "half bond" because, in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which acounts for its paramagnetism and its formal bond order of 2.[4] R-phrases , , , , S-phrases , , , Except where noted otherwise, data are given for materials in their standard state (at 25 Â°C, 100 kPa) Infobox disclaimer and references Nitric oxide or Nitrogen monoxide is a chemical compound with chemical formula NO. This gas is an important signaling molecule in the body of... Simple Illustration of a paramagnetic probe made up from miniature magnets. ...


Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.[4]


Aromatic bond

Main article: Aromaticity

In most cases, the locations of electrons cannot be simplified to simple lines (place for two electrons) or dots (a single electron). In aromatic bonds which occur in planar rings of atoms where the 4n+2 rule determines whether ring molecules would show extra stability. Aromaticity is a chemical property in which a conjugated ring of unsaturated bonds, lone pairs, or empty orbitals exhibit a stabilization stronger than would be expected by the stabilization of conjugation alone. ... The term aromatic compound may also refer to: any organic compound possessing a strong olfactory aroma aromatic hydrocarbons (originally named as a subset of the above; however, aromatic hydrocarbons do not necessarily possess any smell whatsoever) ... In organic chemistry, Hückels rule estimates whether a planar ring molecule will have aromatic properties. ...


In benzene, the prototypical aromatic compound, 18 bonding electrons bind 6 carbon atoms together to form a planar ring structure. The bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.


In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent. Heterocyclic compounds are substances which contain a ring structure as found in benzene and the aromatic compounds, or aromatic hydrocarbons, but in which other atoms than carbon, such as sulfur, oxygen or nitrogen are found as part of the ring. ... For benzine, see petroleum ether. ...


Metallic bond

Main article: Metallic bond

In a metallic bond, 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. Metallic bonds are found in metals like copper. ...


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. In physics, chemistry, and biology, intermolecular forces are forces that act between stable molecules or between functional groups of macromolecules. ... The melting point of a crystalline solid is the temperature range at which it changes state from solid to liquid. ...


Permanent dipole to permanent dipole

Main article: Intermolecular force

A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other. In physics, chemistry, and biology, intermolecular forces are forces that act between stable molecules or between functional groups of macromolecules. ... Electronegativity is a measure of the ability of an atom or molecule to attract electrons in the context of a chemical bond. ... The Earths magnetic field, which is approximately a dipole. ...


Hydrogen bond

Main article: Hydrogen bond

In some ways this is an especially strong example of a permanent dipole, as above. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms, in a three-center two-electron bond like that in diborane. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column. An example of a quadruple hydrogen bond between a self-assembled dimer complex reported by Meijer and coworkers. ... A three-center two-electron bond is an electron deficient chemical bond where three atoms share two electrons. ... The Periodic Table redirects here. ...


Instantaneous dipole to induced dipole (van der Waals)

Main article: van der Waals' forces

Instantaneous dipole to induced dipole, or van der Waals forces, are the weakest, but also the most prolific—occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the (otherwise neutral) atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighbouring helium atom, setting up another dipole. The two atoms will be attracted for an instant, before the charge rebalances and the atoms move on. In chemistry, the term van der Waals forces (sometimes called London dispersion forces) refers to a particular class of intermolecular forces. ... General Name, symbol, number helium, He, 2 Chemical series noble gases Group, period, block 18, 1, s Appearance colorless Standard atomic weight 4. ... Electron cloud is a term used- if not originally coined- by the nobelaurate and acclaimed educator Richard Feynman in The Feynman Lectures on Physics, for discussing exactly what is an electron?. This intuitive model provides a simplified way of visualizing an electron as a solution of the Schrödinger equation. ...


Cation-pi interaction

Main article: Cation-pi interaction

Cation-pi interactions occur between the localized negative charge of π orbital electrons, located above and below the plane of an aromatic ring, and a positive charge. A cation-pi interaction is a noncovalent molecular interaction between the electron-rich orbitals of an aromatic ring with adjacent cation. ... Electron atomic and molecular orbitals, showing a Pi-bond at the bottom right of the picture. ... Properties The electron (also called negatron, commonly represented as e−) is a subatomic particle. ... An aromatic hydrocarbon (abbreviated as AH) or arene [1] is a hydrocarbon, the molecular structure of which incorporates one or more planar sets of six carbon atoms that are connected by delocalised electrons numbering the same as if they consisted of alternating single and double covalent bonds. ...


Electrons in chemical bonds

Many simple compounds involve covalent bonds. These molecules have structures that can be predicted using valence bond theory, and the properties of atoms involved can be understood using concepts such as oxidation number. Other compounds that involve ionic structures can be understood using theories from classical physics. In chemistry, valence bond theory explains the nature of a chemical bond in a molecule in terms of atomic valencies. ... The oxidation number of an element in a molecule or complex is the charge that it would have if all the ligands (basically, atoms that donate electrons) were removed along with the electron pairs that were shared with the central atom[1]. It means that the oxidation number is the... Classical physics is physics based on principles developed before the rise of quantum theory, usually including the special theory of relativity and general theory of relativity. ...


In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. Each atom is assigned an overall electric charge to help conceptualize the molecular orbital's distribution. The forces between atoms (or ions) are largely characterized by isotropic continuum electrostatic potentials. Sodium and chlorine bonding ionically to form sodium chloride. ... Isotropic means independent of direction. Isotropic radiation has the same intensity regardless of the direction of measurement, and an isotropic field exerts the same action regardless of how the test particle is oriented. ...


By contrast, in covalent bonding, the electron density within a bond is not assigned to individual atoms, but is instead delocalized in the MOs between atoms. The widely accepted theory of the linear combination of atomic orbitals (LCAO) helps describe the molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, too, such as Sigma and Pi bond. Covalent redirects here. ... This article may be too technical for most readers to understand. ... This article is being considered for deletion in accordance with Wikipedias deletion policy. ... Electron atomic and molecular orbitals, showing among others the sigma bond of two s-orbitals and a sigma bond of two p-orbitals In chemistry, sigma bonds (σ bonds) are a type of covalent chemical bond. ... Electron atomic and molecular orbitals, showing a Pi-bond at the bottom right of the picture. ...


Atoms can also form bonds that are intermediates between ionic and covalent. This is because these definitions are based on the extent of electron delocalization. Electrons can be partially delocalized between atoms, but spend more time around one atom than another. This type of bond is often called polar covalent. See electronegativity. A polar covalent bond is a form of covalent bonding that happens when atoms of two different elements with different electronegativities bond resulting in an unequal sharing of electrons. ... Electronegativity is a measure of the ability of an atom or molecule to attract electrons in the context of a chemical bond. ...


Thus, the electrons in a molecular orbital (or 'in a polar covalent, or in a covalent bond') can be said to be either localized on certain atom(s) or delocalized between two or more atoms. The type of bond between two atoms is defined by how much the electron density is localized or delocalized among the atoms of the bonds In chemistry, a molecular orbital is a region in which an electron may be found in a molecule. ... Electron density is the measure of the probability of an electron being present at a specific location. ...


References

  1. ^ Laidler, K. J. (1993) The World of Physical Chemistry, Oxford University Press, p. 347
  2. ^ (Journal of Chemical Physics, 1, pg 823, 1933)
  3. ^ Weinhold, F.; Landis, C. Valency and bonding, Cambridge, 2005; pp. 96-100.
  4. ^ a b Pauling, L. The Nature of the Chemical Bond. Cornell University Press, 1960.

External links


  Results from FactBites:
 
Chemical bond - Wikipedia, the free encyclopedia (1882 words)
A chemical bond is the physical phenomenon of chemical substances being held together by attraction of atoms to each other through sharing, as well as exchanging, of electrons or electrostatic forces.
Weak chemical bonds are classically explained to be effects of polarity, or the lack of it, of strong bonds.
The type of strong bond depends on the difference in electronegativity and the distribution of the electron path to the atoms that are bonded.
  More results at FactBites »

 
 

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