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A molecular bond which is also called a covalent bond, is a chemical bond which involves the sharing of electron pairs of atoms. These electron pairs of atoms are known as bonding pairs or shared pairs, and the stable balance of the repulsive and attractive forces between the atoms, when they share electrons, is called covalent bonding. For many of the molecules, those who share electrons, the sharing of electrons allows each and every atom to attain the equivalent of a full outer shell, which is corresponding to a stable electronic configuration. In the stream of organic chemistry, covalent bonds are much more common than ionic bonds, and ionic bonds are very less common.
Many kinds of interactions are there in the covalent bonding, including σ-bonding, metal-to-metal bonding, π-bonding, agostic interactions, bent bonds, and three-center two-electron bonds. The term covalent bond was naed in 1939. Learning about the meaning the prefix co- means jointly, associated in action, partnered to a lesser degree, etc. thus a "covalent bond", means that the atoms share "valence", such as is discussed in valence bond theory of covalent bonds.
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The covalent term in regard to bonding was first used in 1919 by scientist Irving Langmuir in a Journal of the American Chemical Society article which was entitled "The Arrangement of Electrons in Atoms and Molecules". Scientist Langmuir wrote that "we shall denote the term covalent as the number of pairs of electrons that a given atom shares with its neighbor atoms.
The covalent bonding idea can be traced several years before 1919 to Gilbert N. In 1916 Lewis, who described the electron pairs sharing between atoms. He introduced the Lewis Electron or notation dot or Lewis dot structure, in which valence electrons (those in the outer shell) are represented as dotted structures around the atomic symbols. Electron pairs which are located between atoms represent covalent bonds. Multiple pairs represent multiple bonds, such as triple and double. An alternative form of representation, which is not shown here, has bond-forming electron pairs represented as solid lines.
Lewis stated that an atom forms enough covalent bonds to form a full outer electron shell or closed shell. In the diagram of methane shown the carbon atom has a valence of four and is, therefore, surrounded by eight electrons by the octet rule, four from the carbon itself and four from the hydrogens bonded. Each hydrogen has a valence of one and is surrounded by two electrons by the duet rule – its one own electron plus one from the carbon. The numbers of electrons which is correspond to full shells in the quantum theory of the atom; the outer shell of a carbon atom is the n = 2 shell, which can hold eight electrons, whereas the only and outer shell of a hydrogen atom is the n = 1 shell, which can hold only two.
Non-polar Covalent Bonds
When electrons are shared equally between two atoms then a type of nonpolar covalent bond is formed. In an atom, the number of electrons shared by the adjacent atoms will be the same.
A Nonpolar is termed as covalent bond because the difference in electronegativity is mostly negligible. Which means there is no separation of charges between the two atoms or both the atoms have similar electronegativity. When atoms that share a polar bond arrange themselves in such a manner where the electric charges tend to cancel each other then this type of bond is also formed .
Between two identical nonmetal atoms or between different atoms a nonpolar covalent bond can occur.
Non-polar covalent compounds are those in which there is no electronegativity difference. Because there is no bond or dipole moment between the atoms of a molecule and no development of charges on the atoms further, so there is no motion of the bond pair of electrons towards the bonded atoms and that there is no change of electronegativity.
Summary of Covalent Bonds
The shape of many large biological molecules are determined by the noncovalent bonds and stabilize complexes composed of two or more different molecules.
In biological systems there are four main types of noncovalent bonds: ionic bonds, hydrogen bonds, hydrophobic bonds and van der Waals interactions. The bond energies for these interactions range from about 1 to 5 kcal/mol.
A hydrogen atom covalently bonded to an electronegative donor atom associates with an acceptor atom whose nonbonding electrons attract the hydrogen in a hydrogen bond. Among water molecules hydrogen bonds are largely responsible for the properties of both liquid water and the crystalline solid form (ice).
The electrostatic attraction between the positive and negative charges of ions results in Ionic bonds. All cations and anions are surrounded by a tightly bound shell of water molecules in an aqueous solution.
Whenever any two atoms approach each other closely the weak and relatively nonspecific van der Waals interactions are created and they result from the attraction between transient dipoles associated with all molecules.