Bond Energy

Bond Energy is also known as the average bond enthalpy or simply bond enthalpy. It is a quantity that offers insight into the chemical bond strength. The IUPAC bond energy definition of the word 'bond energy' can be given as "the average value that is obtained from the bond dissociation enthalpies (in the gaseous phase) of the entire chemical bonds of a specific type, which is given in a chemical compound.

Thus, the bond energy of a chemical bond in a given compound is visualized as the average amount of energy that is required to break one such chemical bond. 

Note: The bond energy of a chemical bond is always directly proportional to that bond's stability. This signifies that the greater the bond energy of a chemical bond between two atoms, the greater the same chemical bond's stability.

Bond Energy Examples

It is essential to note that a chemical bond's bond energy present in a compound is the average value of the entire individual bond dissociation enthalpies of the chemical bonds. For suppose, the bond energy of the carbon-hydrogen bond in a methane molecule (CH4) is equal to the average bond dissociation energies of every individual carbon-hydrogen bond. We can calculate it as follows:

• CH3 + BDE2 → CH2 + H

• CH4 + BDE1 → CH3 + H

• CH + BDE4 → C + H

• CH2 + BDE3 → CH + H

• BE(C-H) = \[\frac{(BDE_{1}+BDE_{_{2}}+BDE_{3}+BDE_{4})}{4}\]

• Where BDE1 indicates the energy needed to break one carbon-hydrogen bond, present in the CH4 molecule,

• BDE2 indicates the energy needed to break one carbon-hydrogen bond, present in the CH3 molecule,

• BDE3 indicates the energy needed to break one carbon-hydrogen bond, present in the CH2 molecule, and

• BDE4 indicates the energy needed to break the only carbon-hydrogen bond present in the CH molecule.

Finally, the term BE(C-H) indicates the carbon-hydrogen bond's bond energy present in the methane molecule.

Thus, the bond energy of the carbon-hydrogen bond that exists in the methane molecule can be visualized as a change in enthalpy (in general, it is denoted by ΔH) associated by breaking off one CH4 molecule into four hydrogen atoms and one carbon atom, and totally divided by four (since there are four carbon-hydrogen bonds as a total in the methane molecule).

Comparison Between the Bond Energy & Bond Dissociation Energy

Bond dissociation energy of a chemical bond (at times, abbreviated to BDE) is defined as the enthalpy change associated with breaking the chemical bond via homolytic cleavage. For example, the bond dissociation energy of an A-B molecule is the amount of energy needed to facilitate the bond's homolytic cleavage, which exists between A and B, further resulting in the formation of two free radicals.

It is essential to note that the bond dissociation energy of a chemical bond is completely dependent on the environmental absolute temperature. Thus, the bond dissociation energy is usually calculated under the standard conditions (where the temperature is equal to 298 Kelvin, roughly). On the other side, the bond energy of a chemical bond present in a compound is the average value of the total bond dissociation enthalpies of the same bond in the molecule.

Example of a Bond Energy & Bond Dissociation Energy of the Hydrogen-Oxygen Bond in a Water molecule

The bond dissociation energy of hydrogen-oxygen bond in a water molecule can be given as:

H2O + BDE → OH + H

Therefore, bond dissociation energy of hydrogen-oxygen bond in a water molecule can be given as the energy needed to split it into an H and OH free radical.

On the other side, the bond energy of the hydrogen-oxygen bond in the water molecule can be given as:

• OH + BDE2 → H + O

• H2O + BDE1 → OH + H

• BE(O-H) = \[\frac{(BDE_{1}+BDE_{_{2}})}{2}\]

Therefore, the hydrogen-oxygen bond's bond energy in the water molecule is given as the amount of energy required to split the total hydrogen-oxygen bonds in the water molecule, divided by two totally.

Factors Affecting Ionic Bond Energy

Many factors affect the ionic bond energy. An important one among them is given below.

Electronegativity of two atoms bonding together affects the ionic bond energy. In general, the farther away from the electronegativity of 2 atoms, the stronger the bond. 

As an example, Fluorine has the highest, and Cesium has the lowest. They make the strongest ionic bond (at least a well single bond), assuming the Carbon-Fluorine bond is the strongest polar covalent. And mostly, the ionic bonds are stronger than that of the covalent bonds. When checked at melting points, covalent compounds have low melting points, and the ionic compounds have high melting points.