In chemistry, bond energy (E), also called the mean bond enthalpy or average bond enthalpy is the measure of bond strength in a chemical bond. IUPAC defines bond energy as the average value of the gas-phase bond dissociation energies (usually at a temperature of 298 K) for all bonds of the same type within the same chemical species.
The bond dissociation energy (enthalpy) is also referred to as bond disruption energy, bond energy, bond strength, or binding energy (abbreviation: BDE, BE, or D). It is defined as the standard enthalpy change of the following fission: R−X → R + X. The BDE, denoted by Dº(R−X), is usually derived by the thermochemical equation, Dº(R−X) = ΔfHº(R) + ΔfHº(X) – ΔfHº(RX). The enthalpy of formation ΔfHº of a large number of atoms, free radicals, ions, clusters and compounds is available from the websites of NIST, NASA, CODATA, and IUPAC. Most authors prefer to use the BDE values at 298.15 K. 
For example, the carbon–hydrogen bond energy in methane H(C–H) is the enthalpy change (∆H) of breaking one molecule of methane into a carbon atom and four hydrogen radicals, divided by four. The exact value for a certain pair of bonded elemnts varies somewhat depending on the specific molecule, so tabulated bond energies are generally averages from a number of selected typical chemical species containing that type of bond.
Bond energy (E) is the average of all bond-dissociation energies of a single type of bond in a given molecule. The bond-dissociation energies of several different bonds of the same type can vary even within a single molecule. For example, a water molecule is composed of two O–H bonds bonded as H–O–H. The bond energy for H2O is the average of energy required to break each of the two O–H bonds in sequence:
Although the two bonds are the equivalent in the original symmetric molecule, the bond-dissociation energy of an oxygen–hydrogen bond varies depending on whether or not there is another hydrogen atom bonded the oxygen atom.
Standard Bond EnergiesEdit
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Energy is always required to break a bond. Energy is released when a bond is made. In general, the shorter the bond length, the greater the bond energy.
ΔH°* is energy requiring to break a bond (in this article is Kcal/mole).
|Single Bonds||ΔH°* (kJ/mol)||Single Bonds||ΔH°*(kJ/mol)||Multiple Bonds||ΔH°*(kJ/mol)|
Bond energy–distance correlationEdit
Bond strength (energy) can be directly related to the bond length and bond distance. Therefore, we can use the metallic radius, ionic radius, or covalent radius of each atom in a molecule to determine the bond strength. For example, the covalent radius of boron is estimated at 83.0 pm, but the bond length of B–B in B2Cl4 is 175 pm, a significantly larger value. This would indicate that the bond between the two boron atoms is a rather weak single bond. In another example, the metallic radius of rhenium is 137.5 pm, with a Re–Re bond length of 224 pm in the compound Re2Cl8. From this data, we can conclude that the bond is a very strong bond or a quadruple bond. This method of determination is most useful for covalently bonded compounds.
Factors affecting ionic bond energyEdit
Super-Strong, Super-Modulus MaterialsEdit
Melting Temperature–Bond Energy Relation
The bond energy is directly related to the melting temperature of solids.It is seen that for different types of bonds, the melting temperature scales with the bond energy. Both ionic (e.g. NaCl, MgO) and covalent bonds (e.g. Si, C) have high bond energies and consequently high melting temperature. The metallic bond has a large variation in bond energies, ranging from 68 kJ/mol for Hg to 850 kJ/mol for W, matching that of ionic bonds. The lowest bond energies and melting temperatures are for materials with secondary bonds such as NH3, and H2O. Polymers, in which the chains are interconnected by secondary bonds, have very low melting/transition temperatures. The low melting/transition temperatures are related to the breaking of the Van der Waals bond between molecules rather than the covalent bond between atoms.
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