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The term bond-dissociation energy is similar to the related notion of bond-dissociation enthalpy (or bond enthalpy), which is sometimes used interchangeably.However, some authors make the distinction that the bond-dissociation energy (D 0) refers to the enthalpy change at 0 K, while the term bond-dissociation enthalpy is used for the enthalpy change at 298 K (unambiguously denoted DH° 298).
The strength of a bond can be estimated by comparing the atomic radii of the atoms that form the bond to the length of bond itself. For example, the atomic radius of boron is estimated at 85 pm, [10] while the length of the B–B bond in B 2 Cl 4 is 175 pm. [11] Dividing the length of this bond by the sum of each boron atom's radius gives a ratio of
R.T. Sanderson has also noted the relationship between Mulliken electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume. [17] With a knowledge of bond lengths, Sanderson's model allows the estimation of bond energies in a wide range of compounds. [18]
By approximation the bond distance between two different atoms is the sum of the individual covalent radii (these are given in the chemical element articles for each element). As a general trend, bond distances decrease across the row in the periodic table and increase down a group. This trend is identical to that of the atomic radius.
Bond energy and bond-dissociation energy are measures of the binding energy between the atoms in a chemical bond. It is the energy required to disassemble a molecule into its constituent atoms. This energy appears as chemical energy , such as that released in chemical explosions , the burning of chemical fuel and biological processes.
The covalent radius, r cov, is a measure of the size of an atom that forms part of one covalent bond. It is usually measured either in picometres (pm) or angstroms (Å), with 1 Å = 100 pm. In principle, the sum of the two covalent radii should equal the covalent bond length between two atoms, R(AB) = r(A) + r(B).
For a given cation, Pauling defined [2] the electrostatic bond strength to each coordinated anion as =, where z is the cation charge and ν is the cation coordination number. A stable ionic structure is arranged to preserve local electroneutrality , so that the sum of the strengths of the electrostatic bonds to an anion equals the charge on ...
The strong bonding of metals in liquid form demonstrates that the energy of a metallic bond is not highly dependent on the direction of the bond; this lack of bond directionality is a direct consequence of electron delocalization, and is best understood in contrast to the directional bonding of covalent bonds.