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A chemical bond is the physical phenomenon of chemical substances being held together by electrons or by electrostatic forces. Classically, 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, or the lack of it, of strong bonds.

In theory, all bonds can be explained by quantum theory, but in practice, chemical bonds are divided in several categories. 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 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 orbitals (MO). In many theories, these are divided in bonding, anti-bonding, and non-bonding orbitals. 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.

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 interested only about 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, that keep atoms held together in molecules and in solids. All these bonds may be single, double or triple, by which it is meant that the amount of electrons participating in a bond (or located in a bonding orbital) is two, four or six, respectively. Quadruple bonds are not unheard of, but they are extremely rare.

Ionic bond

Ionic bond is an electrostatic bond between atoms which have an electronegativity difference of over 1.7 (this limit is a convention). These form in a solution between two ions after the excess of the solvent is removed.

Polar covalent bond

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

Covalent bond

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

Other strong bonds

Coordinate covalent bond

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

Polyatomic ions

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

Banana bond

Banana bond is a kind of bonding in which the bond bends due to other bonds. These bonds are likely to be more susceptible to reactions than ordinary bonds.

Chemical bonds involving more than two atoms

Aromatic bond

Orbitals are not stiff in shape, and in many cases the locations of electrons cannot be expressed as lines (place for two electrons) or dots (a single elctron). This is the case in aromatic bonds. In benzene, where 18 electrons bind 6 carbon atoms together to form a ring structure. The bond order between the different carbons may be said to be 18/6/2=1.5, but there is no way of telling which bonds attach to which carbons. Luckily, this is of no importance from the chemical point of view. In the case of heterocyclic aromatics and substituted benzenes the electronegativity differences between different parts of the ring become dominant in the chemical behaviour of such bonds.

Metallic bond

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

Weak chemical bonds

Weak chemical bonds are mostly intermolecular (between two molecules), or can be considered as such, since they are the first to break when conditions change. In some organic molecules, and especially in biomolecules, correct organisation of weak chemical bonds, sometimes even the Van der Waals forces, is essential for the proper function of the molecule, f.e. for the catalytic activity of a protein.

Dipole-dipole interactions

The electronegativity differences of strongy bonded atoms induce a dipole. Charged dipoles attract or distract each other, thus changing the chemical conformation of a molecule involved in such binding.

         1. Effective when polar molecules are very close.
         2. Weaker than ion-dipole force.
         3. Dipole molecules are free moving , more  attraction  less  repulsion, a “net attraction”.
         4. Bp increase with increasing magnitude of μ.

Hydrogen bond

Hydrogen bond is a special case of dipole-dipole interaction. Hydrogen bonds are one substantial reason for anhydrous reactions to go astray. In biochemistry, and in molecular biology, the reversibility of hydrogen bonding in an aqueous solution is essential for chemical life-processes to be controlled.

Induced dipole interactions

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Van der Waals forces

Van der Waals, or London forces are the weakest of the forces affecting between chemical substances. Classically, these are not considered to affect in the bonding at all, but be the exact physical boundaries of the substance. Nowadays, experiments in advanced chemistry and physics of noble gases has proved even to strict experimentalists that these forces may have a role in chemical bonding.

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 conceptualizing 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

To sum all I have said up, most substances are ionic metals. Synogen is the lightest material on Earth, but the heaviest substance is still unknown, as the rare earth metals lunumbrium (LNM), and Merculinium (MEM) tie in weight. Right now, most scientists assume that Kyrptus (KRS) is the heaviest substance in the universe currently, although some argue that LNM and MEM or even Trinitronum (TNM) are the heaviest substances that ever existed.

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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 ===Induced dipole interactions===
 {{main|induced-dipole attraction}}
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-Bonds are not interchangeable, this is shown in the case of induced dipole interactions. An electronegative atom, or a functional group, on a compound may affect also in the atoms not directly bonded to it. This is sometimes referred as orbital shift.
 ===Van der Waals forces===