There is a chemical bond when an electrostatic force holds together multiple atoms in a chemical species (strong bonds, or primary or intramolecular) or more molecules in a substance in the condensed state (weak, or secondary, or intermolecular bonds).
The nature of the chemical bond can be explained by observing the coulombic forces interacting between the molecules. Take for example the H2+ cation: it consists of two hydrogen nuclei H and an electron. We denote by Ha the first hydrogen nucleus and with Hb the other hydrogen nucleus. Each of the two nuclei is associated with an electronic wave function, respectively 1sa and 1sb, whose linear combination forms the molecular orbital Ψ.
The molecular orbital Ψ will have low values between the two nuclei, while it will grow closer to them and then decrease further away from them. Therefore if we consider an electron, that is a negative charge placed between the two nuclei, it will be subjected to attractive forces by the two nuclei that will be counterbalanced by the repulsive ones until the stability of the system has been reached; then the electron will be fallen into a potential well from which he will be difficult to escape. In this way, a chemical bond was formed.
Primary chemical bonds
The primary chemical bonds are the forces that hold together the atoms that form the molecules. A primary bond is implemented by the sharing or transfer of electrons between atoms and by the electrostatic attraction between protons and electrons. These bonds generate the transfer of an integer number of electrons, called the bond order, even if in some systems there are intermediate quantities of charge, as in benzene, in which the binding order is 1.5 for each carbon atom. Primary bonds are generally classified into three classes, in order of increasing polarity:
- Covalent bond
- Delocalized bonds and metallic bonding
- Ionic bonding
Covalent bonds are formed between two atoms when both have similar tendencies to attract electrons to themselves (i.e., when both atoms have identical or fairly similar ionization energies and electron affinities).
Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. Because the attraction between molecules, which are electrically neutral, is weaker than that between electrically charged ions, covalent compounds generally have much lower melting and boiling points than ionic compounds.
In fact, many covalent compounds are liquids or gases at room temperature, and, in their solid states, they are typically much softer than ionic solids.
Furthermore, whereas ionic compounds are good conductors of electricity when dissolved in water, most covalent compounds are insoluble in water; since they are electrically neutral, they are poor conductors of electricity in any state.
Delocalized bonds and metallic bonding
Metallic bonding is a type of chemical bonding that rises from the attractive electrostatic force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a structure of positively charged ions (cations).
Ionic bonding results from the electrostatic attraction of oppositely charged ions that are typically produced by the transfer of electrons between metallic and nonmetallic atoms.
Compounds composed of ions are called ionic compounds (or salts), and their constituent ions are held together by ionic bonds: electrostatic forces of attraction between oppositely charged cations and anions. The properties of ionic compounds shed some light on the nature of ionic bonds.
Ionic solids exhibit a crystalline structure and tend to be rigid and brittle; they also tend to have high melting and boiling points, which suggests that ionic bonds are very strong. Ionic solids are also poor conductors of electricity for the same reason—the strength of ionic bonds prevents ions from moving freely in the solid state.
Most ionic solids, however, dissolve readily in water. Once dissolved or melted, ionic compounds are excellent conductors of electricity and heat because the ions can move about freely.
The Formation of ionic compounds
Binary ionic compounds are composed of just two elements: a metal (which forms the cations) and a nonmetal (which forms the anions). For example, NaCl is a binary ionic compound. We can think about the formation of such compounds in terms of the periodic properties of the elements. Many metallic elements have relatively low ionization potentials and lose electrons easily.
These elements lie to the left in a period or near the bottom of a group on the periodic table. Nonmetal atoms have relatively high electron affinities and thus readily gain electrons lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-right corner of the periodic table.
As all substances must be electrically neutral, the total number of positive charges on the cations of an ionic compound must equal the total number of negative charges on its anions. The formula of an ionic compound represents the simplest ratio of the numbers of ions necessary to give identical numbers of positive and negative charges.
Secondary chemical bonds
Secondary chemical bonds are those of the molecular dipoles, which can create the intermolecular attractive forces. The intermolecular bonds are essentially constituted by the mutual attraction between static dipoles (in the case of polar molecules) or between dipoles and ions (this is the case, for example, of a salt which dissolves in water).
In the case of noble gases or compounds formed by apolar molecules the ability to liquefy is explained by the random formation of a temporary dipole when the electrons, in their orbit, are randomly concentrated on one side of the molecule; this dipole induces in the molecules close to itself an imbalance of electric charge (the so-called induced dipole) which generates mutual attraction and causes the gas condensation. The bond is then produced by these particular attraction forces called dispersion or Van der Waals forces.
A particular case of intermolecular bonding, which can also be intramolecular when the geometry of the molecule allows it, is the hydrogen bond.
A hydrogen atom bound to an oxygen (or fluorine) atom, due to its positive polarization and its small size, attracts at relatively high intensity the atoms of oxygen (and fluorine and, to a lesser extent, of nitrogen) neighbors.
This bond, although weak, is responsible for the spatial conformation of proteins and nucleic acids, the conformation on which the biological activity of the compounds depends.
Antibonding molecular orbital
An antibonding molecular orbital is a type of chemical bond given by the overlap of two half-full molecular orbitals. This kind of bond 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.
A molecular-bound orbital becomes an anti-bonded orbital when the electron density between the two nuclei is less than what would be if the two nuclei were separated. Antibonding orbitals are labeled with the asterisk (*) in molecular orbital diagrams; they derive from the out-of-phase superposition of wave functions and are characterized by greater energy than the bound orbitals.
In the case of the diatomic hydrogen (H2) molecule, each atom contributes to the orbital with a single electron, therefore only the orbital σ is occupied and the molecule is more stable than the two separate atoms that compose it.