chemical bond formed when two atoms share one or more pairs of electrons
Molecule


out of two or more elements. Covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are termed "shared pairs" or "bonding pairs", and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed "covalent bonding". Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, and is the primary interaction occurring in ionic compounds. The ions are atoms that have lost one or more
19776552
Chemical bond


shared electrons forms a single covalent bond. The electron density of these two bonding electrons in the region between the two atoms increases from the density of two non-interacting H atoms. A double bond has two shared 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. An example is nitrogen. Quadruple and higher bonds are very rare and
6232674
Chemical bond


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. A single bond between two atoms corresponds to the sharing of one pair of electrons. The Hydrogen (H) atom has one valence electron. Two Hydrogen atoms can then form a molecule, held together by the shared pair of electrons. Each H atom now has the noble gas electron configuration of helium (He). The pair of
6232673
Single bond


In chemistry, a single bond is a chemical bond between two atoms involving two valence electrons. That is, the atoms share one pair of electrons where the bond forms. Therefore, a single bond is a type of covalent bond. When shared, each of the two electrons involved is no longer in the sole possession of the orbital in which it originated. Rather, both of the two electrons spend time in either of the orbitals which overlap in the bonding process. As a Lewis structure, a single bond is denoted
20571378
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. In organic chemistry covalent bonds are much more common than ionic bonds.
6416714
Lewis acids and bases


chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond, it was called a covalent bond. When both electrons come from one of the atoms, it was called a dative covalent bond or coordinate bond. The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen the ammonia molecule donates a pair of electrons to the proton; the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis
5408314
Chemical bond


occur only between certain transition metal atoms. 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 (BF) and ammonia (NH) form an adduct or coordination complex FB←NH with a B–N bond in which a lone pair of electrons on N is shared with an empty atomic orbital on B. BF with an empty orbital is described as an electron pair acceptor or Lewis acid, while NH with a
6232675
History of chemistry


bond is a pair of electrons shared by two atoms. Lewis's model equated the classical chemical bond with the sharing of a pair of electrons between the two bonded atoms. Lewis introduced the "electron dot diagrams" in this paper to symbolize the electronic structures of atoms and molecules. Now known as Lewis structures, they are discussed in virtually every introductory chemistry book. Shortly after the publication of his 1916 paper, Lewis became involved with military research. He did not return to the subject of chemical bonding until 1923, when
1136144
History of molecular theory


lines represent covalent bonds. In this article, he 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' own words: Moreover, he proposed that an atom tended to form an ion by gaining or losing the number of electrons needed to complete a cube. Thus, Lewis structures show each atom in the structure of the molecule using its chemical symbol
12782719
Lewis structure


extend the concept of the electron dot diagram by adding lines between atoms to represent shared pairs in a chemical bond. Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another (pairs of dots can be used instead of lines). Excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms. Although main group elements of the second period and beyond usually react by gaining,
3086896
Valence electron


In chemistry, a valence electron is an outer shell electron that is associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. The presence of valence electrons can determine the element's chemical properties, such as its valence—whether it may bond with other elements and, if so, how readily and with how many. For a main
11942303
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
11727745
20th century in science


Thomson expanded on the work of Wien by showing that charged subatomic particles can be separated by their mass-to-charge ratio, a technique known as mass spectrometry. In 1916, Gilbert N. Lewis published his seminal article "The Atom of the Molecule", which suggested that a chemical bond is a pair of electrons shared by two atoms. Lewis's model equated the classical chemical bond with the sharing of a pair of electrons between the two bonded atoms. Lewis introduced the "electron dot diagrams" in this paper to
22004412
Lone pair


In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom and is sometimes called an unshared pair or non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone pair electrons plus the number of bonding electrons equals the total number of valence electrons around an atom.
9862087
Cubical atom


in the sharing of two electrons. Ionic bonds are formed by the transfer of an electron from one cube to another, without sharing an edge (A). An intermediate state B where only one corner is shared was also postulated by Lewis. Double bonds are formed by sharing a face between two cubic atoms. This results in sharing four electrons: Triple bonds could not be accounted for by the cubical atom model, because there is no way of having two cubes share three parallel edges. Lewis suggested that the electron pairs
3520050
Alfred Lauck Parson


Lewis read a paper by Parson, which argued that the electron, in the Bohr model, might be a ring of negative electricity spinning with a high velocity about its axis and that a chemical bond results from two electrons being shared between two atoms. Parson published the final draft of his theory in 1915. Stimulated by this paper, Lewis published his famous 1916 article "The Atom and the Molecule", in which a chemical bond forms owing to the sharing of pairs of electrons. Several other physicists of the time,
5704348
Skeletal formula


information." A protecting group or protective group is introduced into a molecule by chemical modification of a functional group to obtain chemoselectivity in a subsequent chemical reaction, facilitating multistep organic synthesis. Two atoms can be bonded by sharing more than one pair of electrons. The common bonds to carbon are single, double and triple bonds. Single bonds are most common and are represented by a single, solid line between two atoms in a skeletal formula. Double bonds are denoted by two parallel lines, and triple bonds are shown by three
868179
Chemical bond


lone pair that 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 NH to form the complex ion Ag(NH), which has two Ag←N coordinate covalent bonds. In metallic bonding, bonding electrons are delocalized
6232676
Chemical bond


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 non-polar 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
6232671
Octet rule


bonds by using one s, three p and one d orbitals, or six bonds by using one s, three p and two d orbitals. To form five bonds, the one s, three p and one d orbitals combine to form five spd hybrid orbitals which each share an electron pair with a halogen atom, for a total of 10 shared electrons, two more than the octet rule predicts. Similarly to form six bonds, the six spd hybrid orbitals form six bonds with 12 shared electrons. In this model the
20017342
Chemistry


one electron to become an Na cation while chlorine (Cl), a non-metal, gains this electron to become Cl. The ions are held together due to electrostatic attraction, and that compound sodium chloride (NaCl), or common table salt, is formed. In a covalent bond, one or more pairs of valence electrons are shared by two atoms: the resulting electrically neutral group of bonded atoms is termed a molecule. Atoms will share valence electrons in such a way as to create a noble gas electron configuration
4896873
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. In 1916, G. N. Lewis proposed that a chemical bond forms by the interaction of two shared bonding electrons, with the
19382603
Chemical bond


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 and
6232679
Chemical bond


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
6232656
History of molecular theory


. Moreover, noting that a cube has eight corners Lewis envisioned an atom as having eight sides available for electrons, like the corner of a cube. Subsequently, in 1902 he devised a conception in which cubic atoms can bond on their sides to form cubic-structured molecules. In other words, electron-pair bonds are formed when two atoms share an edge, as in structure C below. This results in the sharing of two electrons. Similarly, charged ionic-bonds are formed by the transfer of an electron from
12782716
Cubical atom


of the periodic table. Although the cubical model of the atom was soon abandoned in favor of the quantum mechanical model based on the Schrödinger equation, and is therefore now principally of historical interest, it represented an important step towards the understanding of the chemical bond. The 1916 article by Lewis also introduced the concept of the electron pair in the covalent bond, the octet rule, and the now-called Lewis structure. Single covalent bonds are formed when two atoms share an edge, as in structure C below. This results
3520049
2-Norbornyl cation


). Non-classical ions differ from traditional cations in their electronic structure: though chemical bonds are typically depicted as the sharing of electrons between two atoms, stable non-classical ions can contain three or more atoms that share a single pair of electrons. In 1939, Thomas Nevell and others attempted to elucidate the mechanism for transforming camphene hydrochloride into isobornyl chloride. In one of the proposed reaction mechanisms depicted in the paper, the positive charge of an intermediate cation was not assigned to a single atom but rather to the structure
2790180
Methanium


intermediate species in chemical reactions. The methanium ion is named after methane (), by analogy with the derivation of ammonium ion () from ammonia (). Methanium can be visualised as a carbenium ion with a molecule of hydrogen interacting with the empty orbital in a 3-center-2-electron bond. The bonding electron pair in the H molecule is shared between the two hydrogen and one carbon atoms making up the 3-center-2-electron bond. The two hydrogen atoms in the H molecule can continuously exchange positions with the three hydrogen atoms in the ion (a
18095860
History of molecular theory


Parson suggested that the electron is not merely an electric charge but is also a small magnet (or "magneton" as he called it) and furthermore that a chemical bond results from two electrons being shared between two atoms. This, according to Lewis, meant that bonding occurred when two electrons formed a shared edge between two complete cubes. On these views, in his famous 1916 article "The Atom and the Molecule", Lewis introduced the "Lewis structure" to represent atoms and molecules, where dots represent electrons and
12782718
Electron pair


two. The pairing of spins is often energetically favorable and electron pairs therefore play a very large role in chemistry. They can form a chemical bond between two atoms, or they can occur as a lone pair of valence electrons. They also fill the core levels of an atom. Because the spins are paired, the magnetic moment of the electrons cancels and the contribution of the pair to the magnetic properties will in general be a diamagnetic one. Although a strong tendency to pair off electrons can be observed in chemistry, it
21865610
Octet rule


to the octet of each bonded atom. In NO, the octet on each atom is completed by four electrons from two two-electron bonds, plus a lone pair of non-bonding electrons on that atom alone. The bond order is 2.5, since each two-electron bond counts as one bond while the three-electron bond has only one shared electron and therefore corresponds to a half-bond. Dioxygen is sometimes represented as obeying the octet rule with a double bond (O=O) containing two pairs of
20017339
Gilbert N. Lewis


formulated the idea of what would become known as the covalent bond, consisting of a shared pair of electrons, and he defined the term odd molecule (the modern term is free radical) when an electron is not shared. He included what became known as Lewis dot structures as well as the cubical atom model. These ideas on chemical bonding were expanded upon by Irving Langmuir and became the inspiration for the studies on the nature of the chemical bond by Linus Pauling. In 1923, he formulated the electron-pair theory of
12768739
Dangling bond


In chemistry, a dangling bond is an unsatisfied valence on an immobilized atom. An atom with a dangling bond is also referred to as an immobilized free radical or an immobilized radical, a reference to its structural and chemical similarity to a free radical. In order to gain enough electrons to fill their valence shells (see also octet rule), many atoms will form covalent bonds with other atoms. In the simplest case, that of a single bond, two atoms each contribute one unpaired electron, and the resulting pair of
18163022
Chemical bond


. Strong chemical bonds are the "intramolecular" forces that 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 bonding is a type of electrostatic interaction between atoms that have a large electronegativity
6232665
Acid


trifluoride (BF), whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly "or" by releasing protons (H) into the solution, which then accept electron pairs. However, hydrogen chloride, acetic acid, and most other Brønsted-
29064
Valence electron


fluorine (F) or chlorine (Cl)). Such an atom has the following electron configuration: sp; this requires only one additional valence electron to form a closed shell. To form an ionic bond, a halogen atom can remove an electron from another atom in order to form an anion (e.g., F, Cl, etc.). To form a covalent bond, one electron from the halogen and one electron from another atom form a shared pair (e.g., in the molecule H–F, the
11942316
Electron


moment of the nucleus is negligible compared with that of the electrons. The magnetic moments of the electrons that occupy the same orbital (so called, paired electrons) cancel each other out. The chemical bond between atoms occurs as a result of electromagnetic interactions, as described by the laws of quantum mechanics. The strongest bonds are formed by the sharing or transfer of electrons between atoms, allowing the formation of molecules. Within a molecule, electrons move under the influence of several nuclei, and occupy molecular orbitals; much as they
9491157
Electron


bond between two atoms is maintained by a pair of electrons shared between them. Later, in 1927, Walter Heitler and Fritz London gave the full explanation of the electron-pair formation and chemical bonding in terms of quantum mechanics. In 1919, the American chemist Irving Langmuir elaborated on the Lewis' static model of the atom and suggested that all electrons were distributed in successive "concentric (nearly) spherical shells, all of equal thickness". In turn, he divided the shells into a number of cells each of which
9491119
Covalent bond


theory. In the molecule , 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. The term "covalence" in regard to bonding was first used in 1919 by Irving Langmuir in a "Journal of the American Chemical Society" article entitled "
6416716
Unpaired electron


In chemistry, an unpaired electron is an electron that occupies an orbital of an atom singly, rather than as part of an electron pair. Each atomic orbital of an atom (specified by the three quantum numbers n, l and m) has a capacity to contain two electrons (electron pair) with opposite spins. As the formation of electron pairs is often energetically favourable, either in the form of a chemical bond or as a lone pair, unpaired electrons are relatively uncommon in chemistry, because an entity that carries an
15135440
Chemical bond


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
6232649
Four-center two-electron bond


A 4-center 2-electron (4c–2e) bond is a type of chemical bond in which four atoms share two electrons in bonding, with a net bond order of . This type of bonding differs from the usual covalent bond, which involves two atoms sharing two electrons (2c–2e bonding). Four-center two-electron bonding is postulated in certain cluster compounds. For instance, the borane anion, is a octahedron with an additional proton attached to one of the triangular faces. As a result, the octahedron is distorted and a B
6900970
Intramolecular force


to chlorine, forming a positively charged sodium ion and a negatively charged chloride ion. In a true covalent bond, the electrons are shared evenly between the two atoms of the bond; there is little or no charge separation. Covalent bonds are generally formed between two nonmetals. There are several types of covalent bonds: in polar covalent bonds, electrons are more likely to be found around one of the two atoms, whereas in nonpolar covalent bonds, electrons are evenly shared. Homonuclear diatomic molecules are purely covalent. The polarity of
3103046
Chemistry


), how atmospheric ozone is formed and how environmental pollutants are degraded (ecology), the properties of the soil on the moon (astrophysics), how medications work (pharmacology), and how to collect DNA evidence at a crime scene (forensics). Chemistry addresses topics such as how atoms and molecules interact via chemical bonds to form new chemical compounds. There are four types of chemical bonds: covalent bonds, in which compounds share one or more electron(s); ionic bonds, in which a compound donates one or
4896848
Alpha cleavage


electrons, 3) sigma bond electrons. One of the lone pair electrons moves down to form a pi bond with an electron from an adjacent (alpha) bond. The other electron from the bond moves to an adjacent atom (not one adjacent to the lone pair atom) creating a radical. This creates a double bond adjacent to the lone pair atom (oxygen is a good example) and breaks/cleaves the bond from which the two electrons were removed. In molecules containing carbonyl groups, alpha-cleavage often competes
7844791
Chemical bond


methane (CH), in which each hydrogen forms a covalent bond with the carbon. See sigma bonds and pi bonds for LCAO descriptions of such bonding. Molecules that 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
6232672
Welding


the molecule as a whole. In both ionic and covalent bonding the location of the ions and electrons are constrained relative to each other, thereby resulting in the bond being characteristically brittle. Metallic bonding can be classified as a type of covalent bonding for which the constituent atoms of the same type and do not combine with one another to form a chemical bond. Atoms will lose an electron(s) forming an array of positive ions. These electrons are shared by the lattice which makes the electron cluster mobile, as the electrons are free
1278838
Three-center two-electron bond


a 3c–2e bond is the trihydrogen cation . This type of bond is also called banana bond. An extended version of the 3c–2e bond model features heavily in cluster compounds described by the polyhedral skeletal electron pair theory, such as boranes and carboranes. These molecules derive their stability from having a completely filled set of bonding molecular orbitals as outlined by Wade's rules. The monomer BH is unstable since the boron atom has an empty p-orbital. A B−H−B 3-center-2-electron bond is formed when a boron atom shares electrons with a B−H bond
11727746
Chemical bonding of H2O


inconsistent with the bond angle of being 104.5°. The actual hybridization of can be explained via the concept of isovalent hybridization or Bent's rule. In short, s character is accumulated in lone pair orbitals because s character is energy lowering relative to p character, and lone pair electrons are closely held with unshared electron density. In contrast, bonding pairs are localized further away and electron density is shared with another atom, so additional s character does not lower energy quite as effectively.  Hence, comparatively more p character is distributed
3364896
Valence (chemistry)


by the sharing of electrons, and ionic bonding leads to octets by the transfer of electrons from one atom to the other. The term covalence is attributed to Irving Langmuir, who stated in 1919 that "the number of pairs of electrons which any given atom shares with the adjacent atoms is called the "covalence" of that atom". The prefix "co-" means "together", so that a co-valent bond means that the atoms share a valence. Subsequent to that, it is now more common to
5443590
Chemical bonding of H2O


and Bent's rule are discussed below to provide a comprehensive bonding model for , explaining and rationalizing the various electronic and physical properties and features manifested by its peculiar bonding arrangements. The Lewis structure of describes the bonds as two sigma bonds between the central oxygen atom and the two peripheral hydrogen atoms with oxygen having two lone pairs of electrons. Valence bond theory suggests that is sp3 hybridized in which the 2s atomic orbital and the three 2p orbitals of oxygen are hybridized to form four new hybridized orbitals which then participate in bonding by overlapping
3364885
Chemical polarity


fourth apex of the approximate tetrahedron, (VSEPR). This orbital is not participating in covalent bonding; it is electron-rich, which results in a powerful dipole across the whole ammonia molecule. In ozone (O) molecules, the two O−O bonds are nonpolar (there is no electronegativity difference between atoms of the same element). However, the distribution of other electrons is unevensince the central atom has to share electrons with two other atoms, but each of the outer atoms has to share electrons with only one other
4893023
Chemical compound


attractive forces that form when the electrons in two adjacent atoms are positioned so that they create a temporary dipole. Additionally, London dispersion forces are responsible for condensing non polar substances to liquids, and to further freeze to a solid state dependent on how low the temperature of the environment is. A covalent bond, also known as a molecular bond, involves the sharing of electrons between two atoms. Primarily, this type of bond occurs between elements that fall close to each other on the periodic table of elements, yet it is
8932400
Covalent bond


orbitals in molecular orbital theory and resonance of sigma bonds in valence bond theory. In three-center two-electron bonds ("3c–2e") 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, with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms.
6416731
Seesaw molecular geometry


AXE in AXE notation). An atom bonded to 5 other atoms (and no lone pairs) forms a trigonal bipyramid; but in this case one of the atoms is replaced by a lone pair. The atom replaced is always an equatorial atom, because the lone pairs repel other electrons more strongly than atoms do. Compounds with disphenoidal geometry (See-Saw Geometry) have two types of ligands: axial and equatorial. The axial pair lie along a common bond axis so that are related by a bond angle of 180
11225839
Chemical bond


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
6232645
Staebler–Wronski effect


It proposes that an electron-hole pair formed by the incident light may recombine near a weak Si–Si bond, releasing energy sufficient to break the bond. A neighbouring H atom then forms a new bond with one of the Si atoms, leaving a dangling bond. These dangling bonds can trap electron-hole pairs, thus reducing the current that can pass through. However, new experimental evidence is casting doubt on this model. More recently, the "H collision model" proposed that two spatially separated recombination events cause
6536206
Coordination complex


atom will give a pair of electrons. There are some donor atoms or groups which can offer more than one pair of electrons. Such are called bidentate (offers two pairs of electrons) or polydentate (offers more than two pairs of electrons). In some cases an atom or a group offers a pair of electrons to two similar or different central metal atoms or acceptors—by division of the electron pair—into a three-center two-electron bond. These are called bridging ligands. Coordination complexes have been known
7538408
Atom


, such as occurs in the compound sodium chloride and other chemical ionic salts. However, many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds. Thus, chemical bonding between these elements takes many forms of electron-sharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds. The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that
336551
Nucleophile


Nucleophile is a chemical species that donates an electron pair 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
16863205
Chemical polarity


is the hydrogen bond, which is also known as the H-bond. For example, water forms H-bonds and has a molar mass M = 18 and a boiling point of +100 °C, compared to nonpolar methane with M = 16 and a boiling point of –161 °C. A molecule may be nonpolar either when there is an equal sharing of electrons between the two atoms of a diatomic molecule or because of the symmetrical arrangement of polar bonds in a more complex molecule. For example, boron trifluoride
4893025
Orbital hybridisation


In chemistry, orbital hybridisation (or hybridization) is the concept of mixing atomic orbitals into new "hybrid orbitals" (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridisation are in fact not
4951798
Aromaticity


the one in the double bond contributes one electron and the one which is not in the double bond and is in a lone pair contributes two electrons to the π system. Polycyclic aromatic hydrocarbons are molecules containing two or more simple aromatic rings fused together by sharing two neighboring carbon atoms (see also simple aromatic rings). Examples are naphthalene, anthracene, and phenanthrene. In fused aromatics, not all carbon–carbon bonds are necessarily equivalent, as the electrons are not delocalized over the entire molecule. The aromaticity of these molecules
9985678
Bonding in solids


typically insulators with large band gaps. The four classes of solids permit six pairwise intermediate forms: Covalent and ionic bonding form a continuum, with ionic character increasing with increasing difference in the electronegativity of the participating atoms. Covalent bonding corresponds to sharing of a pair of electrons between two atoms of essentially equal electronegativity (for example, C–C and C–H bonds in aliphatic hydrocarbons). As bonds become more polar, they become increasingly ionic in character. Metal oxides vary along the iono-covalent spectrum. The Si
14174508
Microstructure


the spatial form of the associated atoms and the molecules and then, at a larger scale, of the crystal, and finally of the organism as a whole. Metallic bonds are formed by the sharing of electrons in the outer layer of the atom in an electron cloud, where they are free and delocalized. This free-electron gas ensures the cohesion of the remaining cations and enables electrical conduction in metals and alloys. Covalent bonds are formed by the sharing of pairs of valence electrons in order to fill the outer electron shells
16529257
Self-healing hydrogels


microstructure, hydrogels are not homogenous within the selected solvent system. The following sections summarize the chemical and physical methods by which hydrogels are crosslinked. Hydrogen bonding is a strong intermolecular force that forms a special type of dipole-dipole attraction. Hydrogen bonds form when a hydrogen atom bonded to a strongly electronegative atom is around another electronegative atom with a lone pair of electrons. Hydrogen bonds are stronger than normal dipole-dipole interactions and dispersion forces but they remain weaker than covalent and ionic bonds. In hydrogels, structure and stability of
7223752
Welding


to another atom to form oppositely charged ions. The bonding in the static position is when the ions occupy an equilibrium position where the resulting force between them is zero. When the ions are exerted in tension force, the inter-ionic spacing increases creating an electrostatic attractive force, while a repulsing force under compressive force between the atomic nuclei is dominant. Covalent bonding takes place when one of the constituent atoms loses one or more electrons, with the other atom gaining the electrons, resulting in an electron cloud that is shared by
1278837
Atom


in its uncombined state is known as the valence shell, and the electrons in that shell are called valence electrons. The number of valence electrons determines the bonding behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills (or empties) their outer valence shells. For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with one-electron more than a filled shell, and others that are one-electron short of a full shell
336550
Charge-shift bond


The charge-shift bond has been proposed as a new class of chemical bond that sits alongside the three familiar families of covalent, ionic bonds, and metallic bonds where electrons are shared or transferred respectively. The charge shift bond derives its stability from the resonance of ionic forms rather than the covalent sharing of electrons which are often depicted as having electron density between the bonded atoms. A feature of the charge shift bond is that the predicted electron density between the bonded atoms is low. It has long been known from experiment that
20070254
History of molecular theory


. Lines are drawn between atoms that are bonded to one another; occasionally, pairs of dots are used instead of lines. Excess electrons that form lone pairs are represented as pair of dots, and are placed next to the atoms on which they reside: To summarize his views on his new bonding model, Lewis states: The following year, in 1917, an unknown American undergraduate chemical engineer named Linus Pauling was learning the Dalton hook-and-eye bonding method at the Oregon Agricultural College, which was the vogue description
12782720
Molecular orbital theory


orbital to an antibonding orbital can occur under UV radiation. This promotion weakens the bond between the two hydrogen atoms and can lead to photodissociation—the breaking of a chemical bond due to the absorption of light. Although in MO theory "some" molecular orbitals may hold electrons that are more localized between specific pairs of molecular atoms, "other" orbitals may hold electrons that are spread more uniformly over the molecule. Thus, overall, bonding is far more delocalized in MO theory, which makes it more applicable to resonant molecules
16289397
Charge-shift bond


is familiar to many, if not all, chemists. The basis of Pauling's description of the chemical bond is that an electron pair bond involves the mixing, resonance, of one covalent and two ionic structures. In bonds between two atoms of the same element, homonuclear bonds, Pauling assumed that the ionic structures make no appreciable contribution to the overall bonding. This assumption followed on from published calculations for the hydrogen molecule in 1933 by Weinbaum and by James and Coolidge. that showed that the contribution of ionic forms amounted to
20070256
Octet rule


shared electrons. However the ground state of this molecule is paramagnetic, indicating the presence of unpaired electrons. Pauling proposed that this molecule actually contains two three-electron bonds and one normal covalent (two-electron) bond. The octet on each atom then consists of two electrons from each three-electron bond, plus the two electrons of the covalent bond, plus one lone pair of non-bonding electrons. The bond order is 1+0.5+0.5=2. Main-group elements in the third and later rows of
20017340
Periodic trends


electron affinity. Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved, using the Pauling scale. Trend-wise, as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down in a group, the electronegativity decreases due to the
9088959
Covalent bond


represented as dots around the atomic symbols. Pairs of electrons located between atoms represent covalent bonds. Multiple pairs represent multiple bonds, such as double bonds and triple bonds. An alternative form of representation, not shown here, has bond-forming electron pairs represented as solid lines. Lewis proposed that an atom forms enough covalent bonds to form a full (or closed) outer electron shell. In the diagram of methane shown here, the carbon atom has a valence of four and is, therefore, surrounded by eight electrons (
6416718
Carbon monoxide


high bond-dissociation energy, the strongest of any neutral molecule, 11.65 eV. Carbon and oxygen together have a total of 10 electrons in the valence shell. Following the octet rule for both carbon and oxygen, the two atoms form a triple bond, with six shared electrons in three bonding molecular orbitals, rather than the usual double bond found in organic carbonyl compounds. Since four of the shared electrons come from the oxygen atom and only two from carbon, one bonding orbital is occupied by two electrons from oxygen, forming a
6359653
Acid


Contrast how the following reactions are described in terms of acid-base chemistry: In the first reaction a fluoride ion, F, gives up an electron pair to boron trifluoride to form the product tetrafluoroborate. Fluoride "loses" a pair of valence electrons because the electrons shared in the B—F bond are located in the region of space between the two atomic nuclei and are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF is a Lewis acid because it accepts the electron pair from
29074
Electron


electrons versus positive nuclei changes the binding energy of an atomic system. The exchange or sharing of the electrons between two or more atoms is the main cause of chemical bonding. In 1838, British natural philosopher Richard Laming first hypothesized the concept of an indivisible quantity of electric charge to explain the chemical properties of atoms. Irish physicist George Johnstone Stoney named this charge 'electron' in 1891, and J. J. Thomson and his team of British physicists identified it as a particle in 1897. Electrons can also participate in nuclear reactions,
9491100
Chemical bond


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 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
6232657
Electron deficiency


electron-deficient molecules are typically strongly electron-attracting (electrophilic). As the most extreme form of electron deficiency one can consider the metallic bond. For many years, "electron-deficiency" was often used as a general descriptor for boron hydrides and other molecules featuring multicenter bonding (in which a pair of bonding electrons extends over more than two atoms, as in three-center two-electron bonds), as a way of distinguishing such molecules from conventionally bonded compounds such as hydrocarbons. However, this usage was
21953956
Octet rule


the periodic table can form hypercoordinate or hypervalent molecules in which the central main-group atom is bonded to more than four other atoms, such as phosphorus pentafluoride, PF, and sulfur hexafluoride, SF. For example, in PF, if it is supposed that there are five true covalent bonds in which five distinct electron pairs are shared, then the phosphorus would be surrounded by 10 valence electrons in violation of the octet rule. In the early days of quantum mechanics, Pauling proposed that third-row atoms can form five
20017341
Bond valence method


same bonding strengths. In practice some tolerance is allowed, but bonds are rarely formed if the ratio of the bonding strengths of the two atoms exceeds two, a condition expressed by the inequality shown in Eq. 4. This is known and the "valence matching rule". Atoms with non-bonding valence electrons, i.e., with lone pairs, have more flexibility in their bonding strength than those without lone pairs depending on whether the lone pairs are stereoactive or not. If the lone pairs are not stereoactive, they are
21896879
William Lipscomb


carboranes (compounds of carbon, boron, and hydrogen). The three-center two-electron bond is illustrated in diborane (diagrams at right). In an ordinary covalent bond a pair of electrons bonds two atoms together, one at either end of the bond, the diboare B-H bonds for example at the left and right in the illustrations. In three-center two-electron bond a pair of electrons bonds three atoms (a boron atom at either end and a hydrogen atom in the middle),
15233060
Gilbert N. Lewis


but Lewis’ version of this theory was the only one to be embodied in a concrete atomic model. Again Lewis’ theory did not interest his Harvard mentors, who, like most American chemists of that time, had no taste for such speculation. Lewis did not publish his theory of the cubic atom, but in 1916 it became an important part of his theory of the shared electron pair bond. In 1916, he published his classic paper on chemical bonding ""The Atom and the Molecule"" in which he
12768738
Chemistry


Shell Electron Pair Repulsion model (VSEPR), and the concept of oxidation number can be used to explain molecular structure and composition. An ionic bond is formed when a metal loses one or more of its electrons, becoming a positively charged cation, and the electrons are then gained by the non-metal atom, becoming a negatively charged anion. The two oppositely charged ions attract one another, and the ionic bond is the electrostatic force of attraction between them. For example, sodium (Na), a metal, loses
4896872
Cyanate


The cyanate ion is the anion with the chemical formula written as [OCN] or [NCO]. In aqueous solution it acts as a base, forming isocyanic acid, HNCO. The cyanate ion is an ambidentate ligand, forming complexes with a metal ion in which either the nitrogen or oxygen atom may be the electron-pair donor. It can also act as a bridging ligand. Organic cyanates are called isocyanates when there is a C−NCO bond and cyanates when there is a C−OCN bond. The three atoms in a cyanate
22019152
Molecular geometry


geometry is determined by the quantum mechanical behavior of the electrons. Using the valence bond approximation this can be understood by the type of bonds between the atoms that make up the molecule. When atoms interact to form a chemical bond, the atomic orbitals of each atom are said to combine in a process called orbital hybridisation. The two most common types of bonds are sigma bonds (usually formed by hybrid orbitals) and pi bonds (formed by unhybridized p orbitals for atoms of main group elements). The geometry can also be
13558558
Technetium


halides form anionic clusters with Tc–Tc bonds. The situation is similar for the related elements of Mo, W, Re. These clusters have the nuclearity Tc, Tc, Tc, and Tc. The more stable Tc and Tc clusters have prism shapes where vertical pairs of Tc atoms are connected by triple bonds and the planar atoms by single bonds. Every technetium atom makes six bonds, and the remaining valence electrons can be saturated by one axial and two bridging ligand halogen atoms such as chlorine or bromine. Technetium forms
10740691
Valence (chemistry)


" and can form only one bond with an atom that donates a "valence electron" to complete chlorine's outer shell. However, chlorine can also have oxidation states from +1 to +7 and can form more than one bond by donating "valence electrons". Hydrogen has only one valence electron, but it can form bonds with more than one atom. In the bifluoride ion (), for example, it forms a three-center four-electron bond with two fluoride atoms: Another example is the Three-center
5443601
Covalent bond


"n" = 1 shell, which can hold only two. While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, quantum mechanics is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. Walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond (molecular hydrogen) in 1927. Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the
6416720
Antibonding molecular orbital


In chemical bonding theory, an antibonding orbital is a type of molecular orbital (MO) that weakens the chemical bond between two atoms and helps to raise the energy of the molecule relative to the separated atoms. Such an orbital has one or more nodes in the bonding region between the nuclei. The density of the electrons in the orbital is concentrated outside the bonding region and acts to pull one nucleus away from the other and tends to cause mutual repulsion between the two atoms. This is in contrast to a bonding molecular orbital
698921
Intramolecular force


main types of chemical bonds — ionic, covalent, and metallic — distinguished by the degree of charge separation between participating atoms. The characteristics of the bond formed can be predicted by the properties of constituent atoms, namely electronegativity. They differ in the magnitude of their bond enthalpies, a measure of bond strength, and thus affect the physical and chemical properties of compounds in different ways. % of ionic character is directly proportional difference in electronegitivity of bonded atom. An ionic bond involves a complete transfer of one or more valence electrons
3103044
Antibonding molecular orbital


, which has a lower energy than that of the separate atoms, and is responsible for chemical bonds. Antibonding molecular orbitals (MOs) are normally "higher" in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. If two hydrogen atoms are initially far apart, they have identical atomic orbitals. However, as the spacing between the two atoms becomes smaller, the electron wave functions begin to overlap. The Pauli exclusion principle prohibits any two electrons (e-) in a molecule from having
698922
Pi bond


In chemistry, pi bonds (π bonds) are covalent chemical bonds where two lobes of an orbital on one atom overlap two lobes of an orbital on another atom and this overlap occurs laterally. Each of these atomic orbitals has zero electron density at a shared nodal plane, passing through the two bonded nuclei. The same plane is also a nodal plane for the molecular orbital of the pi bond. The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as
16289193
Ether


presence of two lone pairs of electrons on the oxygen atoms makes hydrogen bonding with water molecules possible. Cyclic ethers such as tetrahydrofuran and 1,4-dioxane are miscible in water because of the more exposed oxygen atom for hydrogen bonding as compared to linear aliphatic ethers. Other properties are: Ethers are quite stable chemical compounds which do not react with bases, active metals, dilute acids, oxidising agents, and reducing agents. Generally, they are of low chemical reactivity, but they are more reactive than alkanes. Epoxides, ketals, and acetals
9276689
Chemical compound


. The metals in ionic bonding usually lose their valence electrons, becoming a positively charged cation. The nonmetal will gain the electrons from the metal, making the nonmetal a negatively charged anion. As outlined, ionic bonds occur between an electron donor, usually a metal, and an electron acceptor, which tends to be a nonmetal. Hydrogen bonding occurs when a hydrogen atom bonded to an electronegative atom forms an electrostatic connection with another electronegative atom through interacting dipoles or charges. A compound can be converted to a different chemical composition by
8932402
Octet rule


bond are counted twice, once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, two (shown in red) from the oxygen itself and two (shown in black) from the carbon. All four of these electrons are counted in both the carbon octet and the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity (such as sodium) and the second a nonmetal of high electronegativity (such as chlorine).
20017327
Ammonia


theory (VSEPR theory) with an experimentally determined bond angle of 106.7°. The central nitrogen atom has five outer electrons with an additional electron from each hydrogen atom. This gives a total of eight electrons, or four electron pairs that are arranged tetrahedrally. Three of these electron pairs are used as bond pairs, which leaves one lone pair of electrons. The lone pair of electrons repel more strongly than bond pairs, therefore the bond angle is not 109.5°, as expected for a regular tetrahedral arrangement, but 106.7°
1341323
Digermyne


the two Ge atoms bears partial positive charge and is electron deficient, the other Ge atom has an electron lone pair and is able to donate some electron density. The abnormal bond angles and the single-bond feature of LGeGeL can be rationalized by the electron donating character of N atom, which leads to the formation of N p(π)→Ge (empty p orbital) interaction. Therefore, the donor-acceptor bonds between two Ge atoms are weakened and are more like non-bonding electron lone pairs. It has been suggested that the
1702571
Spin chemistry


fashions (T, T,T). Spin states relate to chemical and biochemical reaction mechanisms because bonds can be formed only between two electrons of opposite spin (Hund’s Rule). Sometimes when a bond is broken in a particular manner, for example, when struck by photons, each electron in the bond relocates to each respective molecule, and a radical-pair is formed. Furthermore, the spin of each electron previously involved in the bond is conserved, which means that the radical-pair now formed is a
10873123