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Types of chemical bond


What is a chemical bond ?

The chemical bond is a concept that refers to the forces of attraction that manifest at the molecular and atomic level . It is usually distinguished between intramolecular chemical bonds and intermolecular chemical bonds:

  • Chemical bond intramolecular : are the forces of attraction between atoms and ions that hold them together to form the molecules. They occur with electron exchange between the bonded atoms. For example, the covalent bond and the ionic bond . They are stronger than intermolecular bonds.
  • Intermolecular chemical bonds : are the forces of attraction between molecules . They are mainly due to electrostatic forces. For example, the hydrogen bridges or the dispersion forces of London.

Models on the chemical bond

The concept of chemical bonding is very complex and there are several theories that try to explain its formation and characteristics. It is very common to talk about the octet rule to explain why chemical bonding occurs. According to this rule, the formation of the chemical bond is produced following a basic principle of chemistry according to which every system tends to evolve towards a situation of lower energy and more stable. This situation in the atoms would occur when each atom has 8 electrons in its external orbital, just like the noble gases, elements whose atoms are so stable that they do not react with each other and isolated atoms can be found existing naturally.

Although the octet rule can help predict the formation of a chemical bond, it is actually a theoretical approximation or simplification, as is also the Valence Layer electron pairs repulsion theory (TREPEV).

The qualities of the bond and the effects on matter are usually explained through electrostatics while the formation of intramolecular bonds are explained mainly by the valence bond theory and the theory of molecular orbitals, none of which has been given definitively.

Valence bond theory

This theory was formulated in 1927 by Heitler and London based on the idea proposed by Lewis in 1916 according to which a chemical bond is formed by the joint interaction of two electrons shared by two atomic nuclei . Heitler and London proposed the idea that the electrons involved in the bonds are the valence electrons , these are the electrons of the last energy level. There are elements that can exist with a different number of valence electrons, which is explained by different oxidation states , because at the end of the day the oxidation number would represent the same as the atomic valence.

The chemical bond would be a process in which two atoms share valence electrons to complete their last level of energy and thus arrive at a more stable situation . The electrons are shared by overlap of atomic orbitals, and the hybridization and resonance of these orbitals can occur. The noble gases have a valence of zero because they have their last complete energy level and hence their practically null reactivity and are the only elements that exist naturally as isolated atoms.

The so-called modern valence bond theory replaces the overlap of atomic orbitals with the overlap of valence bond orbitals, orbitals that expand throughout the molecule and is generally seen as a theory complementary to the theory of molecular orbitals; each can explain and predict more directly some properties and characteristics of chemical bonds. For example, the valence bond theory predicts the behavior of homonuclear diatomic molecules (H 2 , O 2 , F 2 , etc.) much more accurately than the theory of molecular orbitals.

The theory of molecular orbitals

In the theory of molecular orbitals the overlap of orbitals between two atoms is replaced by their linear combination to form orbitals that cover the entire molecule and in which the electrons participating in the bond would be located. According to this theory, as many molecular orbitals are formed as atomic orbitals are combined and each one includes several nuclei, usually only 2.

The theory of molecular orbitals better predicts the spectroscopic, magnetic and ionization behavior of molecules in general. It also best describes pentavalent molecules, metal bonding and electron-deficient systems.


Currently the two theories are considered approximations of a better theory, the interaction of complete configurations . This theory combines all the possible electronic configurations of the system from linear formality; Whether it is applied to the valence bond theory or applied to the theory of molecular orbitals, the same wave function is reached, hence the two theories are considered approximations of this wave function. It is considered that the valence bond theory is a localized link approach and that the theory of molecular orbitals is a delocalised approach.

Types of chemical bond

The chemical bond can be classified into two major groups, the intramolecular chemical bonds and the intermolecular chemical bonds, each with several subtypes.

Intramolecular chemical bonds

These bonds are what keep the atoms together inside a molecule . The three main types are the covalent bond, the ionic bond and the metal bond.

Covalent bond

Covalent bond

Hydrogen-carbon covalent bond

The covalent bond is formed when electrons are shared between two nuclei . Generally an even number of electrons participate, frequently 2, 4 or 6, because the molecular energy is usually lower when in an orbital there are two electrons paired with opposite spin. The covalent bond can be simple, double, triple, quadruple, … according to 2, 4, 6, … electrons. None of the atoms yields electrons to the other atom, the net charge of each atom and of the molecule as a whole is neutral .

If the electrons are attracted equally by the two nuclei, the charges are distributed evenly in the molecule. However, when one of the atoms presents greater electronegativity , it attracts the electrons of the link more towards it, diverting the center of the charge; the result is an electrically neutral molecule but with a non-symmetric charge distribution that is measured by the dipole moment.

Covalent bonds between atoms with a difference in electronegativity greater than 0.4 are usually referred to as polar covalent bonds because they produce an inhomogeneous distribution of charges. In fact, all covalent bonds are polar, except if they occur between identical atoms that, having a similar electronegativity, form an apolar covalent bond with a homogeneous charge distribution.

A special type of covalent bond is the dative or coordinated covalent bond , in which the pair of electrons participating in the bond comes from a single atom. Other types of covalent bond are the aromatic bond, the bonds of one and three electrons, the flexed bond or the deficient links of electrons (link of 2 centers and 3 electrons (2c-3e) and link of 3 centers and 4 electrons (3c -4e)).

Ionic bond

Ionic bond
Ionic bond NaF (sodium – fluorine)

The ionic bond is due to electrostatic attraction between two atoms that have a large difference in electronegativity, such as two ions of opposite charge . An exact value of electronegativity difference that produces an ionic bond is not known, but a difference greater than 2 is usually cited as the limit from which ionic bonds are formed; below 1.7, polar covalent bonds would form.

The characteristics of the ionic bond can be explained through the formation of an anion (negative charge ion) and a cation (positive charge ion) or as if it were a covalent bond between two atoms with a difference in electronegativity so high that atom yields an electron to the most electronegative atom.

Metallic link

The metallic bond is produced when a set of atoms share unlocalized electrons among all of them. If the metal bond is compared with the ionic bond, in the ionic bond the location of the charges is static whereas in the metal bond it is not, which is responsible for the electrical conductivity of many metals.

Vibrational link

The vibrational bond is a type of bond that was proposed at the beginning of the 1980s and that occurs when a very light atom vibrates rapidly between two others of much higher mass . Its existence has been much discussed without it being accepted yet. Recently, in a study published on the 10th of 2014 , we have obtained data that could confirm its existence as observed in the reaction between Muonio and Bromo .

Muonio is an exotic atom that consists of a nucleus formed by a single particle of antimuon (positive charge, is the antiparticle of the muon) orbited by an electron. This makes the muonium even lighter than the protium , the lighter isotope of hydrogen. When the muonium is between two bromine atoms, much heavier, the muonium transcends between the two vibrating rapidly reducing the total energy of the system and joining the bromine atoms (BrMuBr). This peculiar link would explain why the reactivity of this molecule decreases with increasing temperature. More studies on the vibrational link are necessary to corroborate its existence and check if it could also be given among other different elements.

Intermolecular chemical bonds

Intermolecular bonds are formed by forces of attraction between two molecules . The four main types are dipole-dipole and its types, the hydrogen bond, the dispersion forces of London and the cation-pi interaction. These links are responsible for many of the characteristics of the substances, for example the melting point.


When two atoms are joined with a large difference in electronegativity, the charge of the molecule is distributed asymmetrically. The electrons participating in the bond will be closer to the more electronegative atom, thus producing a partial negative charge of the molecule in this area and a positive partial charge in the opposite zone. This molecule with neutral net charge but asymmetric charge distribution is known as a dipole . Two or more molecules of this type can interact with each other by electrostatic forces; for example, two dipoles will be attracted when approaching the areas with opposite partial load. These electrostatic forces can also occur between a dipole and other charged particles, for example ions.

A type of dipole-dipole interaction is the interaction between a permanent dipole and what is known as an induced dipole . Some molecules do not have dipolar characteristics by themselves but the electrostatic force exerted by a dipole (permanent dipole) can generate an asymmetric charge distribution in the other molecule, forming the induced dipole.

Hydrogen bond

The one known as hydrogen bond or bridge is actually a strong dipole-dipole interaction . It occurs in molecules formed by the union of hydrogen and oxygen, nitrogen or fluorine. Between hydrogen and any of these elements there is a large difference in electronegativity that causes a high dipole moment responsible for strong electrostatic interactions between molecules. For example, hydrogen bonding is responsible for the high boiling point of water relative to heavier analog molecules since more energy is needed to overcome the hydrogen bond and separate the water molecules to go into vapor state.

London DISPERSION forces

As the distribution of the charge of the electrons is not uniform at all times with respect to the nucleus, there is always some asymmetric distribution of charges around an atom forming instantaneous dipoles that can induce a dipole in neighboring molecules, hence the dispersion forces of London are also known as instantaneous dipole-induced dipole forces . This type of interactions are very weak compared to other intermolecular interactions and are manifested even in apolar molecules. The larger the size of the molecule, the dispersion forces in London become more intense as instantaneous dipoles can be formed more easily. They are considered a type of forces of var der Waals.

Cation-Pi interaction

Cation-Pi interaction between benzene and sodium
Cation-Pi interaction between benzene and sodium

The interaction cation-pi (can be seen as cation-π) is the interaction that appears between molecules rich in electrons located in pi orbitals and a cation (positive charge ion). For example, benzene, and in general aromatic hydrocarbons , are electron-rich systems in pi orbitals that can interact with positive ions such as Li + or Na +. The intensity of these interactions can become similar to a hydrogen bond and, like hydrogen bonding, they have a very important role in nature, for example the cation-PI interaction is very important in the structure of proteins and in the enzymatic activity.

  • Weinhold, F. and Landis, C. (2005). Valency and bonding . Cambridge. pp. 96-100. ISBN 0-521-83128-8.
  • W. Locke (1997). Introduction to Molecular Orbital Theory . Retrieved May 18, 2005.
  • Emilio San Fabián Maroto (01-March-2012). Interaction of Configurations . Department of Physical Chemistry, University of Alicante.
  • AJ Stone (1996), The Theory of Intermolecular Forces , Oxford: Clarendon Press.
  • Donald G. Fleming 1, Jörn Manz2,3, Kazuma Sato and Toshiyuki Takayanagi (December 8, 2014.). “Fundamental Change in the Nature of Chemical Bonding by Isotopic Substitution”. Angewandte Chemie International Edition; 53 (50): 13706-13709. doi: 10.1002 / anie.201408211 .
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