Bond type, bond length and bond energy. Electronegativity.

Chemical Bonding

Atoms or ions in a compound or an elementary substance are held together by electromagnetic attraction forces called chemical bonds. Bond type, bond length and bond energyareimportant characteristicsof a particular chemical bond.

There are three main bond types: ionic, covalent and metallic.Additionally, molecules of a substance are always attracted by weak intermolecular forces(Van der Waals bond). The kind of chemical bond depends on the electronic structure of atoms forming this bond and ability of these atoms to gain or lose electrons.

Covalent bondsare formed by sharing of electrons between atoms;commonly electrons of two adjacent atoms are paired. The electron density of atoms entering the covalent bond is shifted to the region between the nuclei. The binding force results from the electromagnetic attraction between the shared electrons and the positive nuclei of the atoms. Covalent bonds are formed by atoms of two nonmetallic elements or atoms of nonmetallic and metallic elements, when the last is in a high oxidation state. Valence is a number of covalent bonds formed by an atom; this term is used to describe an atom’s bond-forming capacity.

Ionic bondsare formed by transfer of electrons from one atom to another. Electrostatic interaction held the ions appeared together. Typical metals react with active nonmetals to form ionic compounds. Oxidation states (oxidation numbers) are positive or negative integer numbers assigned to atoms in a compound supposed that all bonds are completely ionic. The oxidation numbers are useful for balancing red-ox reactions; but they are close to real charges on atoms for ionic compounds only.

Metallic bondsare formed by delocalization of bonding electrons over a lattice composed of ions. Only the electrons of the valent shell are delocalized, the electrons of the inner shell are bonded with individual nuclei of the metallic element.

Bond length is a distance between centers of bonded atoms or ions and bond energy (orbond strength) is quantity of energy, necessary to destroy the bond. The strength of chemical bonds varies considerably. Chemical properties of substances depend on chemical bonds type and strengths.

The energy curve for a diatomic molecule is shown in fig. 3.1. Minimum of energy in this curve is observed at interatomic distance equal to the bond length. The depth of this minimum equals to the bond energy.

Fig. 3.1. The energy curve for a diatomic molecule.

Electronegativity. When a chemical bond is formed, electronic density is rearranged. Ability of an isolated atom to gain or lose electrons depends on its electron affinity and ionization energy. To characterize an atom in a compound concept of electronegativity was proposed by Linus Pauling. Electronegativity is a measure of the ability of a specified atom in a compound to attract electrons of outer shells. L. Pauling observed that a bond between different atoms is almost always stronger then a bond between the same atoms, due to shift of electron density from one atom to another and attraction between the opposite charges appeared. The extra bond strength of molecule AB (DE) can be calculated as:

DE = E_A-B – ½(E_A-A + E_B-B),

where E_A-A, E_B-B and E_A-B are energies for the A-A, B-B and A-B bonds, respectively.

A difference in electronegativities of the atoms A and B is proportional to the extra bond strength. L. Pauling calculated differences in electronegativity for various pairs of atoms and used these values to develop a scale of electronegativity. The most electronegative element is fluorine and the less electronegative element is francium; electronegativity values for F and Fr are 4.0 and 0.7, respectively.

Besides Pauling’s electronegativity scale there are several other ones. For instance, R.S. Mulliken used the average of the ionization energy and electron affinity to calculate electronegativities of elements. The values of electronegativity of particular atom vary slightly in different scales. The interrelations between values of electronegativity of the atoms forming the bond and bond types are given in the table 3.1.

Table. 3.1. Electronegativities of the atoms and bond types.

* Atom 2 is an atom of the same element as atom 1.

Octet rule:atoms entering chemical bonds often tend to gain, loose or share electrons until there are eight electrons in the outer shell. This tend is caused by an extra stability of filled ns²np⁶ (n³2) outer electron shells[1]. The first elements of the Periodic Table: H, He, Li and Be don’t obey the octet rule. For B atom octet can be completed in some compounds if a coordinate (donor-acceptor) covalent bond is formed (see below). Except for the alkali, alkali-earth metallic elements and aluminum, the octet rule doesn’t work well for cations. Failure of the octet rule is often observed for transition elements or elements like tin or lead, located in the Periodic Table next to transition elements. Some of these elements form pseudo-noble-gas configuration, containing filled d-subshell, for example Cd²⁺ [Ar]4s²4p⁶4d¹⁰.

Lewis symbols and Lewis structures. Chemical bonds between atoms can be represented using Lewis symbols and Lewis structures. To construct a Lewis symbol for an element, we write its atomic symbol surrounded by a number of dots (or ´’s, or circles, etc.), each of them represents one electron in the atom’s valent shell. For example, hydrogen, which has one electron in its valent shell, is given the Lewis symbol H×. Two electrons paired are represented by a pair of dots; number of unpaired electrons in a Lewis symbol is maximal of possible. (Lewis symbols usually are not applied to the transition elements.) The formulas constructed using Lewis symbols are calledLewis structures or electron dot formulas. The Lewis structure for H₂ is H:H, the pair of electrons in the bond is shown by a pair of dots between the two H atoms. Often a dash is used to represent a covalent bond instead of the pair of dots (for example, H–H and :NºN:). The Lewis structures are most useful for describing covalent bonds.