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. History See also Atomic theory, Atomism The concept that matter is composed of discrete units and cannot be divided into arbitrarily tiny In Chemistry, a molecule is defined as a sufficiently stable electrically neutral group of at least two Atoms in a definite arrangement held together by A chemical compound is a substance consisting of two or more different elements chemically bonded together in a fixed proportion by Mass. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. Quantum electrodynamics ( QED) is a relativistic Quantum field theory of 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. Quantum mechanics is the study of mechanical systems whose dimensions are close to the Atomic scale such as Molecules Atoms Electrons 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. In Chemistry, a molecule is defined as a sufficiently stable electrically neutral group of at least two Atoms in a definite arrangement held together by In Materials science, a crystal is a Solid in which the constituent Atoms Molecules or Ions are packed in a regularly ordered repeating Structure is a fundamental and sometimes Intangible notion covering the Recognition, Observation, nature, and Stability of

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". An ionic bond (or electrovalent bond) is a type of Chemical bond that can often form between Metal and Non-metal Ions (or A hydrogen bond results from a Dipole-dipole force between an Electronegative atom and a Hydrogen atom bonded to Nitrogen, Oxygen The Van der Waals equation is an Equation of state that can be derived from a special form of the potential between a pair of molecules (hard-sphere repulsion Care should be taken because the strongest of the "weak" bonds can be stronger than the weakest of the "strong" bonds.

Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis structures, also called Lewis-dot diagrams are diagrams that show the bonding between Atoms of a Carbon (kɑɹbən is a Chemical element with the symbol C and its Atomic number is 6 Hydrogen (ˈhaɪdrədʒən is the Chemical element with Atomic number 1 Oxygen (from the Greek roots ὀξύς (oxys (acid literally "sharp" from the taste of acids and -γενής (-genēs (producer literally begetteris the Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.

Overview

The electrons of atoms are electromagnetically attracted by the nuclei of atoms, due to the opposite electric charge of elecrons and nuclei. Chemical bonds are characterised by physical states in which a few electrons move partly from one atom to one or more other atoms, driven by the achievement of a lower state of energy from this motion. This lowering of energy is caused by a rearrangement of charges, usually resulting in net decrease in the average distance between the electrons of all the bonded atoms, and their nuclei. The transfer of charge caused by the movement of the electron from one atom to another, also causes the participating atoms (which may number from two to many) to be attracted to one another electromagnetically. The attractive force between atoms is the bond.

Chemical bonds, for the sake of simplicity, are classically assigned characteristics of two major types: covalent and ionic.

In a simplified view of a pure covalent type bond, the bond forms as a few electrons farthest from their atomic nuclei become more attracted to the region of space between two nuclei. In this region, negatively-charged electrons experience attraction from the positively-charged protons from more than one nucleus. This causes some electrons to spend a high probability of their time in the interatomic space. In turn, the nuclei are stabilized in position by the pull from these shared electrons during the fraction of time that the bonding electrons reside between the atoms. Although such bonding electrons do not spend all of their time between atoms, when they spend more time between a given pair of atoms than otherwise, they constitute chemical bonds. Nuclei fixed by such bonds may vibrate, but they are pulled toward each other by the mutual forces of the bonding electrons pulling them together, yet prevented from approaching too closely by their own charge, or else by the mutual repulsion of other inner electrons, which are held so closely and tightly to individual nuclei that they cannot be shared to any important degree.

In a simplified view of a pure ionic type bond, one or more outer electrons are not shared between atoms, but donated from one atom to another. In such a bond, the structure of the electron cloud of one of the nuclei contains an available space for another electron, which allows an additional electron to experience a greater net attraction from the nucleus than is experienced by outer electrons in a neighboring atom toward their own nucleus. This difference in available states causes effective transfer of one or more electrons from one atom to another atom, where they can be more tightly bound. This transfer causes the donating atom to assume a net positive charge, and the other to assume a net negative charge; the atoms thus become positive or negatively charged ions. An ion is an Atom or Molecule which has lost or gained one or more Valence electrons giving it a positive or negative electrical charge The bond then results from electrostatic attraction between these ionized atoms.

Most bonds have a mixture of covalent and ionic character, as bonding electrons are shared between atoms, but shared somewhat unevenly. 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. 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 Valence shell electron pair repulsion (VSEPR theory (1957 is a model in Chemistry, which is used for predicting the shapes of individual Molecules based 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. In Chemistry, valence bond theory explains the nature of a Chemical bond in a Molecule in terms of atomic valencies. -->In Chemistry Resonance in Chemistry is a theory used to represent and model certain types of non-classical Molecular structures Resonance is a key component A linear combination of atomic orbitals or LCAO is a Quantum superposition of Atomic orbitals and a technique for calculating Molecular orbitals Ligand field theory (LFT describes the bonding in Coordination complexes. Electrostatics are used to describe bond polarities and the effects they have on chemical substances. Electrostatics is the branch of Science that deals with the Phenomena arising from what seems to be stationary Electric charges Since Classical

History

Main articles: History of chemistry and History of the molecule

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. The history of Chemistry begins with the discovery of Fire, then Metallurgy which allowed purification of metals and the making of alloys as well as the exploitation In Chemistry, the history of the Molecule traces the origins of the concept or idea of the existence of strong chemical bonds between two or more Atoms Chemical species are Atoms Molecules molecular fragments Ions etc In Chemical physics and Physical chemistry, chemical affinity can be defined as electronic properties by which dissimilar Chemical species are capable of 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". Sir Isaac Newton, FRS (ˈnjuːtən 4 January 1643 31 March 1727) Biography Early years See also Isaac Newton's early life and achievements Opticks is a book written by English physicist Isaac Newton that was released to the public in 1704. History See also Atomic theory, Atomism The concept that matter is composed of discrete units and cannot be divided into arbitrarily tiny In Physics, a force is whatever can cause an object with Mass to Accelerate. 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:

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. A voltaic pile is a set of individual Voltaic cells placed in series Friherre Jöns Jacob Berzelius (20 August 1779 &ndash 7 August 1848 was a Swedish chemist By the mid 19th century, Edward Frankland, F. Sir Edward Frankland, KCB, FRS ( January 18, 1825 &ndash August 9, 1899) was a Chemist, one of the foremost 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. Adolph Wilhelm Hermann Kolbe ( September 27, 1818 &ndash November 25, 1884) was a German Chemist. In Chemistry, radicals (often referred to as free radicals) are atoms molecules or ions with Unpaired electrons on an otherwise Open shell In Chemistry, valence, also known as valency or valency number, is a measure of the number of Chemical bonds formed by the Atoms 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' own words:

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. Walther Ludwig Julius Kossel ( January 4, 1888 in Berlin, Germany &ndash 22 May, 1956 in Tübingen, Germany was a German Both Lewis and Kossel structured their bonding models on that of Abegg's rule (1904). 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. Walter Heinrich Heitler ( 2 January 1904 &ndash 15 November 1981) was a German physicist who made contributions to Quantum electrodynamics Fritz Wolfgang London ( March 7, 1900 &ndash March 30, 1954) was a German -born American theoretical Physicist. The Heitler-London method forms the basis of what is now called valence bond theory. In Chemistry, valence bond theory explains the nature of a Chemical bond in a Molecule in terms of atomic valencies. 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. A linear combination of atomic orbitals or LCAO is a Quantum superposition of Atomic orbitals and a technique for calculating Molecular orbitals Sir John Edward Lennard-Jones KBE FRS (born October 27, 1894; died November 1, 1954) was a Mathematician who held a chair of theoretical Fluorine, fluorum meaning "to flow" is the Chemical element with the symbol F and Atomic number 9 Oxygen (from the Greek roots ὀξύς (oxys (acid literally "sharp" from the taste of acids and -γενής (-genēs (producer literally begetteris the 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. In Chemistry, a molecular orbital (or MO) is a region in which an Electron may be found in a Molecule. In Physics, especially Quantum mechanics, the Schrödinger equation is an equation that describes how the Quantum state of a Physical system 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. Quantum chemistry is a branch of Theoretical chemistry, which applies Quantum mechanics and Quantum field theory to address issues and problems in Density functional theory (DFT is a quantum mechanical theory used in Physics and Chemistry to investigate the Electronic structure (principally

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 Chemistry, valence bond theory explains the nature of a Chemical bond in a Molecule in terms of atomic valencies. In chemistry valence electrons are the Electrons contained in the outermost or valence, Electron shell of an Atom. An atomic orbital is a Mathematical function that describes the wave-like behavior of an electron in an atom 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”. Linus Carl Pauling (February 28 1901 – August 19 1994 was an American Scientist, Peace activist, Author and educator. 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:

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

Main article: Molecular orbital theory

Molecular orbital theory (MO) uses a linear combination of atomic orbitals to form molecular orbitals which cover the whole molecule. In Chemistry, molecular orbital theory ( MO theory) is a method for determining molecular structure in which Electrons are not assigned to individual In Chemistry, a molecular orbital (or MO) is a region in which an Electron may be found in a Molecule. An atomic orbital is a Mathematical function that describes the wave-like behavior of an electron in an atom These are often divided into bonding orbitals, anti-bonding orbitals, and non-bonding orbitals. Antibonding (or anti-bonding) is a type of chemical bonding. An antibonding orbital is a form of Molecular orbital (MO that is located outside the region A molecular orbital is merely a Schrödinger orbital which includes several, but often only two nuclei. In Chemistry, a molecular orbital (or MO) is a region in which an Electron may be found in a Molecule. 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. Antibonding (or anti-bonding) is a type of chemical bonding. An antibonding orbital is a form of Molecular orbital (MO that is located outside the region 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. An atomic orbital is a Mathematical function that describes the wave-like behavior of an electron in an atom These electrons neither contribute nor detract from bond strength.

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 Chemistry, bond energy ( E) is a measure of Bond strength in a Chemical bond. 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.

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. Friedrich Hund (4 February 1896 - 31 March 1997 was a German Physicist from Karlsruhe known for his work on atoms and molecules Robert Sanderson Mulliken ( June 7, 1896 &ndash October 31, 1986) was an American physicist and chemist Gerhard Herzberg, PC, CC, FRSC, FRS ( December 25, 1904 &ndash March 3, 1999) was a pioneering 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. The term complex in Chemistry is usually used to describe molecules or ensembles formed by the combination of Ligands and metal Ions. Electron deficiency occurs when a compound has too few Valence electrons for the connections between atoms to be described as Covalent bonds Electron deficient bonds g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made. Diborane is the Chemical compound with the formula B2H6 It is a colorless gas at room temperature with a repulsively 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. Quantum chemistry is a branch of Theoretical chemistry, which applies Quantum mechanics and Quantum field theory to address issues and problems in However better valence bond programs are now available.

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. A chemical formula is a way of expressing information about the Atoms that constitute a particular Chemical compound, and how the relationship between those atoms changes Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Organic chemistry is a discipline within Chemistry which involves the scientific study of the structure properties composition reactions, and preparation In Organic chemistry, functional groups are specific groups of Atoms within Molecules that are responsible for the characteristic Chemical reactions 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. In Chemistry, conformational isomerism is a form of Stereoisomerism in which Molecules with the same Structural formula (same connectivity 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.

Strong chemical bonds

These chemical bonds are intramolecular forces, which hold atoms together in molecules. Hydrogen (ˈhaɪdrədʒən is the Chemical element with Atomic number 1 Carbon (kɑɹbən is a Chemical element with the symbol C and its Atomic number is 6 Nitrogen (ˈnaɪtɹəʤɪn is a Chemical element that has the symbol N and Atomic number 7 and Atomic weight 14 Oxygen (from the Greek roots ὀξύς (oxys (acid literally "sharp" from the taste of acids and -γενής (-genēs (producer literally begetteris the Abundance Owing to their high Reactivity, the halogens are found in the environment only in compounds or as Ions Halide ions and oxoanions Sulfur or sulphur (ˈsʌlfɚ see spelling below) is the Chemical element that has the Atomic number 16 In Chemistry, a molecule is defined as a sufficiently stable electrically neutral group of at least two Atoms in a definite arrangement held together by 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. In Chemistry, bond strength is measured between two Atoms joined in a Chemical bond. For example, the carbons in benzene are connected to each other with about 1. Benzene, or benzol, is an organic Chemical compound and a known Carcinogen with the molecular formula C 6 H 6 5 bonds, and the two atoms in nitric oxide NO, are connected with about 2. Nitric oxide or nitrogen monoxide is a Chemical compound with Chemical formula N[[Oxygen O]] 5 bonds. Quadruple bonds are also well known. A quadruple bond is a type of Chemical bond between two Atoms involving 8 Electrons This bond is an extension of the more familiar types Double bonds 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. " Electronegativity " is the opposite of " Electropositivity," which describes an element's ability to donate electrons 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.

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. An organic compound is any member of a large class of Chemical compounds whose Molecules contain Carbon. See sigma bonds and pi bonds for LCAO-description of such bonding. In Chemistry, sigma bonds ( σ bonds) are the strongest type of covalent Chemical bond. In Chemistry, pi bonds ( π bonds) are covalent Chemical bonds where two lobes of one involved electron orbital overlap two lobes

Polar covalent bond

Main article: Polar covalent bond

Polar covalent bonding is intermediate in character between a covalent and an ionic bond. "Polar molecule" and "Non-polar" redirect here

Ionic bond

Main article: Ionic bond

Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. An ionic bond (or electrovalent bond) is a type of Chemical bond that can often form between Metal and Non-metal Ions (or There is no precise value that distinguishes ionic from covalent bonding but a difference of electronegativity of over 2. 0 is likely to be ionic and a difference of less than 1. 5 is likely to be covalent. ^[3] Ionic bonding leads to separate positive and negative ions. An ion is an Atom or Molecule which has lost or gained one or more Valence electrons giving it a positive or negative electrical charge Ionic charges are commonly between −3e to +3e. The elementary charge, usually denoted e, is the Electric charge carried by a single Proton, or equivalently the negative of the electric charge carried The elementary charge, usually denoted e, is the Electric charge carried by a single Proton, or equivalently the negative of the electric charge carried

Coordinate covalent bond

Main article: 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. A coordinate covalent bond (formerly also known as dative bond) is a description of Covalent bonding between two atoms in which both electrons shared in the bond come This concept is somewhat fading as chemists increasingly embrace molecular orbital theory. Examples of coordinate covalent bonding occur in nitrones and ammonia borane. A nitrone is the N-oxide of an Imine and a Functional group in Organic chemistry. Ammonia borane (or borazane is the Chemical compound with the formula H3NBH3 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.

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 bond, also known as banana bond, is a term in Organic chemistry that refers to a type of covalent Chemical bond with a geometry somewhat Bent bond, also known as banana bond, is a term in Organic chemistry that refers to a type of covalent Chemical bond with a geometry somewhat See also Intramolecular forces ' Steric effects arise from the fact that each Atom within a Molecule occupies a certain Bent bonds are often more susceptible to reactions than ordinary bonds.

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

In three-center two-electron bonds three atoms share two electrons in bonding. A three-center two-electron bond is an electron deficient Chemical bond where three atoms share two electrons This type of bonding occurs in electron deficient compounds like diborane. Diborane is the Chemical compound with the formula B2H6 It is a colorless gas at room temperature with a repulsively sweet odor 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. The 3-center-4-electron bond is a model used to explain bonding in Hypervalent molecules such as Phosphorus pentafluoride, Sulfur hexafluoride, the 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 In certain cluster compounds so-called four-center two-electron bonds also have been postulated. 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 These bonds are not to be confused with "two-center four-electron bonds", which is one way of describing conventional double bonds between two atoms. The MO description of such double bonds would distinguish one bond, with electron density concentrated along the line between atoms, as a σ bond, but the other bond, with electron density concentrated in lobes on either side of the axial line, as a π (pi) bond.

One- and three-electron bonds

Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. In Chemistry, radicals (often referred to as free radicals) are atoms molecules or ions with Unpaired electrons on an otherwise Open shell 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. Dilithium, Li2 is a Diatomic molecule comprising two Lithium Atoms covalently bonded together This exception can be explained in terms of hybridization and inner-shell effects. ^[4]

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. Nitric oxide or nitrogen monoxide is a Chemical compound with Chemical formula N[[Oxygen O]] The oxygen molecule, O₂ can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2. Paramagnetism is a form of magnetism which occurs only in the presence of an externally applied magnetic field ^[5]

Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities. ^[5]

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. In Organic chemistry, Hückel's 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. Benzene, or benzol, is an organic Chemical compound and a known Carcinogen with the molecular formula C 6 H 6

Metallic bond

Main article: Metallic bond

In a metallic bond, bonding electrons are delocalized over a lattice of atoms. Although the term metallic bond is often used in contrast to the term Covalent bond it is better to speak of metallic bonding, because this type of bonding is By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. Because of delocalization or the free moving of electrons, it leads to the metallic properties such as conductivity, ductility and hardness.

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. In Physics, Chemistry, and Biology, intermolecular forces are forces that act between stable Molecules or between functional groups of Often, these define some of the physical characteristics (such as the melting point) of a substance. The melting point of a 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). In Physics, Chemistry, and Biology, intermolecular forces are forces that act between stable Molecules or between functional groups of " Electronegativity " is the opposite of " Electropositivity," which describes an element's ability to donate electrons In physics there are two kinds of dipoles ( Hellènic: di(s- = two- and pòla = pivot hinge An electric dipole is a Dipoles will attract or repel each other.

Hydrogen bond

Main article: Hydrogen bond

In some ways this is an especially strong example of a permanent dipole, as above. A hydrogen bond results from a Dipole-dipole force between an Electronegative atom and a Hydrogen atom bonded to Nitrogen, Oxygen 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. A three-center two-electron bond is an electron deficient Chemical bond where three atoms share two electrons 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. The periodic table of the chemical elements is a tabular method of displaying the Chemical elements Although precursors to this table exist its invention is

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. The Van der Waals equation is an Equation of state that can be derived from a special form of the potential between a pair of molecules (hard-sphere repulsion 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. Helium ( He) is a colorless odorless tasteless non-toxic Inert Monatomic Chemical Electron cloud is a term used if not originally coined by the Nobel Prize laureate and acclaimed educator Richard Feynman in The Feynman Lectures on Physics 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.

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. Cation-\pi interaction is a noncovalent molecular interaction between the face of an electron-rich \pi system (e In Chemistry, pi bonds ( π bonds) are covalent Chemical bonds where two lobes of one involved electron orbital overlap two lobes The electron is a fundamental Subatomic particle that was identified and assigned the negative charge in 1897 by J

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. In Chemistry, valence bond theory explains the nature of a Chemical bond in a Molecule in terms of atomic valencies. The oxidation number of a central atom in a coordination compound is the charge that it would have if all the Ligands were removed along with the Electron pairs Other compounds that involve ionic structures can be understood using theories from classical physics.

In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. An ionic bond (or electrovalent bond) is a type of Chemical bond that can often form between Metal and Non-metal Ions (or 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. Isotropy is uniformity in all directions Precise definitions depend on the subject area

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. A linear combination of atomic orbitals or LCAO is a Quantum superposition of Atomic orbitals and a technique for calculating Molecular orbitals Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. Anisotropy (pronounced with stress on the third syllable ˌænaɪˈsɒtrəpi is the property of being directionally dependent as opposed to Isotropy, which means homogeneity These may have their own names, too, such as Sigma and Pi bond. In Chemistry, sigma bonds ( σ bonds) are the strongest type of covalent Chemical bond. In Chemistry, pi bonds ( π bonds) are covalent Chemical bonds where two lobes of one involved electron orbital overlap two lobes

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. "Polar molecule" and "Non-polar" redirect here See electronegativity. " Electronegativity " is the opposite of " Electropositivity," which describes an element's ability to donate electrons

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. In Chemistry, a molecular orbital (or MO) is a region in which an Electron may be found in a Molecule. 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

References

External links

• W. Locke (1997). Introduction to Molecular Orbital Theory. Retrieved May 18, 2005. Events 1152 - Henry II of England marries Eleanor of Aquitaine. Year 2005 ( MMV) was a Common year starting on Saturday (link displays full calendar of the Gregorian calendar.
• Carl R. Nave (2005). HyperPhysics. Retrieved May 18, 2005. Events 1152 - Henry II of England marries Eleanor of Aquitaine. Year 2005 ( MMV) was a Common year starting on Saturday (link displays full calendar of the Gregorian calendar.
• Linus Pauling and the Nature of the Chemical Bond: A Documentary History. Retrieved February 29, 2008. Leap years Although the modern calendar counts a year as 365 days a complete revolution around the sun takes approximately 365 days and 6 hours 2008 ( MMVIII) is the current year in accordance with the Gregorian calendar, a Leap year that started on Tuesday of the Common