Chemical bond

A chemical bond is the physical phenomenon of chemical substances being held together by attraction of atoms to each other through sharing, as well as exchanging, of electrons or electrostatic forces. In general, strong chemical bonds are found in molecules, crystals or in solid metal and they organize the atoms in ordered structures. Weak chemical bonds are classically explained to be effects of polarity between molecules which contain strong polar bonds. Some very weak bond-like interactions also result from induced polarity London forces between the electron clouds of atoms, or molecules. Such forces allow the liquification and solidification of inert gases. At the very lowest strengths of such interactions, there is no good operational definition of what constitutes a proper definitional "bond".

Overview

In the simplest view of a so-called covalent bond, an electron or pair of electrons may be drawn into the space between two atomic nuclei, because in this region the negative electrons are subject to the positive charge of both nuclei, instead of just one of them. At the same time, electrons present between nuclei screen the repulsion between nuclei, and cause them to be attracted to the electrons (which are closer), and thus toward each other, instead of repelled by each other. This situation tends to hold the nuclei and electrons in a relatively fixed configuration, although all are still free to move in accordance with the dictates of quantum mechanics.

In a simplified view of an ionic bond, one or more electrons is simply transferred from one atom to another, causing one atom to assume some of the character of a positive ion, and the other a negative ion. The bond then results from electrostatic attraction between atoms. Some of the complexity of the process is hidden in this description, in that the reason why an atom would transfer an electron to another, is a complicated matter which also involves quantum theory.

In theory, all bonds can be explained by quantum theory, but in practice, chemical bonds are divided in several categories, as above. Simplifications of quantum theory have been developed to describe and predict the bonds and their properties. These theories include octet theory, valence bond theory, orbital hybridization theory, VSEPR theory, ligand field theory and LCAO -method. Electrostatics and other physical theories are used to describe bond polarities and the effects they have on chemical substances. Actual chemical bonds are not exactly described by these theories, due to the uncertainty principle. However, in combination, they constitute a powerful theory, which can be applied in almost all of chemistry.

In quantum mechanics, in simplified terms, electrons are located on an atomic orbital (AO), but in a strong chemical bond, they form a molecular orbital (MO), as noted above. In many theories, these are divided into bonding, anti-bonding, and non-bonding orbitals (according to the regions where electrons tend to be found). A molecular orbital is merely a Schrödinger orbital which includes two (and occasionally more) nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being between nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an anti-bonding orbital and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly atomic orbitals) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength.

Molecular orbitals are further divided according the types of atomic orbitals hybridizing to form a bond. These orbitals are results of electron-nucleus interactions that are caused by the fundamental force of electromagnetism. Chemical substances will form a bond if their orbitals become lower in energy when they interact with each other. Different chemical bonds are distinguished that differ by electron cloud shape and by energy levels.

History

Early speculations into the nature of the chemical bond, from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. By the mid 19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called “combining power”, in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert Lewis developed the concept of the electron-pair bond. Walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of molecular hydrogen, in 1927, using a valence bond theory. In 1930, a first mathematically complete quantum description of the simplest chemical bond (that produced by one electron in the hydrogen molecular ion) was derived in the Ph.D. thesis of Edward Teller.

In 1931, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, Heitler, and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were generally known:

^* the electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.

^* the spins of the electrons have to be opposed.

^* once paired, the two electrons can not take part in additional bonds.

His last three rules were new:

^* the electron-exchange terms for the bond involves only one wave function from each atom.

^* the available electrons in the lowest energy level form the strongest bonds.

^* 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 result to become what some have called the “bible” of modern chemistry.

Bonds in chemical formulas

The 3-dimensionality of atoms and molecules makes it hard to use a single technique for indicating orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH₃-CH₂-OH), separating the functional group from another part of the molecule (C₂H₅OH), or by its atomic constituents (C₂H₆O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons ( with the 2-dimensionalized approximate directions) are marked, f.e. for elemental carbon _.^'C_.^' Some chemists may also mark the respective orbitals, f.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 keep atoms held together in molecules and in solids. Quite often these bonds will be single, double or triple in strength, that is, the number of electrons participating in a bond (or located in a bonding orbital) is two, four, or six, respectively. Even numbers are common because electrons enjoy lower energy states, if paired. Substantially more advanced bonding theories have shown that bond strength is not always a whole number, depending on the distribution of electrons to each atom involved in a bond. For example, the carbons in benzene are connected to each other with about 1.5 bonds, and the two atoms in nitric oxide NO, are connected with about 2.5 bonds. Quadruple bonds are not unheard of, but they are extremely rare. The type of strong bond depends on the difference in electronegativity and the distribution of the electron orbital paths available to the atoms that are bonded. The larger the 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 electronegativity, the more covalent properties (full sharing) the bond has.

Covalent bond

Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. In the latter case, the bond is sometimes referred to as purely covalent. See sigma bonds and pi bonds for current LCAO-explanation of non-polar bonds.

Polar covalent bond

Polar covalent bonding is by nature an intermediate type of bond, between a covalent bond and an ionic bond. In more advanced theories of bonding, all bonds may be considered somewhat polar.

Ionic bond

Ionic bonding is a type of electrostatic interaction between atoms which have an electronegativity difference of over 1.6 (this limit is a convention). These form in a solution between two ions after the excess of the solvent is removed. The strongest form of chemical bond is the ionic bond between two ions of opposite charges, and such high energies are responsible for the stability and high melting points of ionic crystals ("salts"). Ionic charges are commonly between -3e to +7e

Other strong bonds

Coordinate covalent bond

Coordinate covalent bonding is a special type of bonding, in which the bonding electrons originate solely from another atom. This is different from an ionic bond in that the electronegativity difference is small.

Polyatomic ions

A different type of bond between two atoms happens commonly in ions. The bond is located in the midst of three (or more) atoms. This occurs usually in polyatomic ions such as methanoate (or formate) (HCOO^-) anion, in which the 0,5 order bond carries the effective charge of -1.

Banana bond

The Banana bond is a kind of bonding in which the bond bends, often due to the presence of an influencing atom in the middle of another covalent bond. These bonds are likely to be more susceptible to reactions than ordinary bonds. An example can be found in diborane, where the bonds between boron atoms are known as "three center, two-electron bonds". 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.

Chemical bonds involving more than two atoms

Aromatic bond

Orbitals often have complex shapes and densities, and in many cases the locations of electrons cannot be simplified to simple lines (place for two electrons) or dots (a single electron). This is the case in aromatic bonds which occur in rings of atoms where the 4n+2 rule determines whether ring molecules comprised of C=C bonds would show behavior extra stability by allowing extra sharing of electrons below and above the ring plane.

In benzene, the protypical 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 carbons 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.

Metallic bond

A metallic bond, as an ionic bond (strictly), exists only in a solid (or liquid) state. In a metallic bond, there are delocalized electrons in a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are quite static.

Intermolecular bonding

There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms.

Permanent dipole to permanent dipole

A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other.

Hydrogen bond

In some ways this is an especially srong example of a permanent dipole bond, as above. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms, in a three-center system like the banana bond in diborane. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column.

Instantaneous dipole to induced dipole (Van der Waals)

Instantaneous dipole to induced dipole, or Van der Waals forces, are the weakest, but also the most prolific - occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the - otherwise-neutral - atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighboring 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

Cation-pi interactions occur between the localized negative charge of $\pi$ orbital electrons, located above and below the plane of an aromatic ring, and a positive charge.

Electrons in chemical bonds

Many simple compounds involve covalent bonds. These molecules have structures that can be predicted using valence bond theory, and the properties of atoms involved can be understood using concepts such as oxidation number. Other compounds that involve ionic structures can be understood using theories from classical physics.

In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. Each atom is assigned an overall electric charge to help conceptualize the molecular orbital's distribution. The forces between atoms (or ions) are largely characterized by isotropic continuum electrostatic potentials.

By contrast, in covalent bonding, the electron density within a bond is not assigned to individual atoms, but is instead delocalized in the MOs between atoms. The widely-accepted theory of the linear combination of atomic orbitals (LCAO) helps describe the molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, too, such as Sigma and Pi bond

Atoms can also form bonds that are intermediates between ionic and covalent. This is because these definitions are based on the extent of electron delocalization. Electrons can be partially delocalized between atoms, but spend more time around one atom than another. This type of bond is often called polar covalent. See electronegativity.

Thus, the electrons in a molecular orbital (or 'in a polar covalent, or in a covalent bond') can be said to be either localized on certain atom(s) or delocalized between two or more atoms. The type of bond between two atoms is defined by how much the electron density is localized or delocalized among the atoms of the substance.

Limitations of valence bond theory

However, more complicated compounds such as metal complexes, or electron deficient compounds, cannot be described by valence bond theory alone, and quantum chemistry (based on quantum mechanics) has to be used.

Linus Pauling's book The Nature of the Chemical Bond has influenced the development of chemistry concerning bond formation as the increasingly complex theories are required.

Determination of chemical properties through chemical bonding

Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics, such as the melting point) of a substance. These forces include ionic interactions, hydrogen bonds, dipole-dipole interactions, and induced dipole interactions.

References

W. Locke (1997). Introduction to Molecular Orbital Theory. Retrieved May 18, 2005.
Carl R. Nave (2005). HyperPhysics. Retrieved May 18, 2005.

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