Bond Enthalpy: Definition, Average Bond Enthalpies, and Enthalpy of a Reaction

Bond enthalpy (also known as bond-dissociation enthalpy, average bond energy, or bond strength) is the amount of energy contained in a bond formed between atoms in a molecule.

• It refers to the amount of energy required for a bond to be cleaved in the gas phase, whether homolytic or symmetrically.
• A homolytic or symmetrical bond-breaking event occurs when the bond is broken.
• Each atom that originally participated in the bond gains one electron, transforming into a radical rather than an ion.

Key Terms: Enthalpy, Bond Enthalpy, Covalent Bond, Bond strength, Endothermic reaction, Exothermic reaction, Diatomic molecules, Bond-dissociation enthalpy, Average Bond Enthalpy

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What is Bond Enthalpy?

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Bond enthalpy, also known as bond energy, is a quantity that provides insight into the strength of a chemical bond and thus its stability. The bond enthalpy of a chemical bond is the total amount of energy necessary to break one mole of that bond.

• For example, the bond enthalpy of the oxygen-hydrogen single bond is 463 kJ per mole.
• This means that 463 kilojoules of energy are required to break one mole of hydrogen-oxygen single bonds. 
• It is essential to remember that breaking a chemical bond is always an endothermic process.
• Because the molecule requires energy to break the chemical bonds that make it up.
• Breaking a chemical bond always results in a positive enthalpy change (ΔH > 0).
• In contrast, the formation of a chemical bond is often an endothermic process.
• In such cases, the enthalpy change will be negative (ΔH < 0).
• The term 'mean bond enthalpy' or 'average bond enthalpy' is used to represent the strength of a single, particular bond in a molecule.
• The mean bond energy can be determined by calculating the average values of all the bond dissociation energies in the molecule.
• Thus, mean bond enthalpy differs from bond dissociation energy (with the exception of diatomic molecules). 
• Bond enthalpy can also be stated as the amount of energy required to break one mole of a chemical bond at 298K in the gas phase.
• The larger the value, the stronger the bond and the more energy necessary to break it.

The image below shows how the chemical bond between atoms A and B is broken when energy equal to the bond enthalpy is supplied to molecule AB.

Bond Enthalpy

For example, the bond dissociation energy required to break down one mole of gaseous hydrogen chloride molecule into gaseous hydrogen and chlorine atoms is 432kJ, while the bond dissociation enthalpy of gaseous HCl is +432kJ per mol.

• If a molecule contains multiple bonds, the bond enthalpy is calculated for each bond, and the average value is used.
• For example, methane (CH₄) has four C-H bonds with an average bond energy of +1652 kJ and +415.5 kJ per mole of bond.

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Average Bond Enthalpies

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Bond enthalpy gives us an insight into the strength of the bond, therefore it is an important parameter to be considered. In order to express the strength of a single bond the 'average bond enthalpy' or 'mean bond enthalpy’ is noted.

• It can be calculated by taking the average of all the average bond enthalpies of that particular type of bond present in a given molecule.
• Hence, the mean bond enthalpy is different from the average bond enthalpy.

Several studies have already resulted in the formation of a table indicating the average bond enthalpy of several single and multiple bonds as illustrated in the figure below.

But in actuality, the bond enthalpy is also affected by the presence of adjacent bonds, and the experimental values are not always equal to the theoretical values.

• For example in the case of methane, as we know the bond enthalpy for individual CH bonds is 413 kJ/mol.
• Therefore theoretically the bond enthalpy for all the CH bonds in methane, four bonds, should have been 413 X 4 = 1659 kJ.
• But instead, it is 1660 kJ/mol, having a variation of 1kJ/mol.

Moreover, the bond enthalpy values also depend on the state of matter, in cases of gases the bond enthalpy is closer to its average value, but in the case of liquids, it is higher since a greater amount of energy is required to convert the liquid into a gaseous state, and highest in case of solids.

Hence, the factors affecting bond enthalpy are:

• Electronegativity of the atoms that are making the bond
• Nature and state of matter have the bonds of liquid, gaseous, and solids.

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Enthalpy of the Reaction

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As we know enthalpy refers to the amount of energy absorbed or released from a system under constant pressure, and the change in enthalpy due to any chemical reaction is termed as enthalpy of reaction. It is represented by ΔH_rxn.

Mathematically, the enthalpy of the reaction can be calculated using the following formulas:

ΔH_rxn = Potential energy of product bonds – Potential energy of reactant bonds

⇒ ΔH_rxn = energy added to break the bonds (+ ΔH) + energy released while forming the bonds (- ΔH)

One approach to calculating the enthalpy of a reaction is to individually calculate the bond enthalpies of all the bonds present in the molecule.

Endothermic and Exothermic Process

At the time of the reaction, energy is absorbed to break the bond and this is called endothermic energy.

• When energy is released to form a new bond is known as the exothermic process.
• In endothermic reactions, the products are in higher energy than reactants, and the energy difference (ΔH) between them is always positive.
• In exothermic reactions, the reactants are in higher energy than the products, and the energy difference between them.
• This is called a change in enthalpy of the reaction and is always negative.

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Estimation of Enthalpy of a Reaction 

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The enthalpy of reaction can be calculated by following the five steps mentioned below:

• Step 1: Identification of the bonds that would break during the reaction and determine their bond enthalpies (+ ΔH).
• Step 2: Take the sum of the bond enthalpies for the bonds broken.
• Step 3: Identification of the bonds that would form during the reaction and determine their bond enthalpies (- ΔH).
• Step 4: Take the sum of the bond enthalpies for the products formed.
• Step 4: Combine both the values obtained in Step 2 and Step 4. We would obtain the ΔH_rxn.

Example: Hydrogenation of propene to propane in the presence of hydrogen gas. Here propene is reduced in the presence of hydrogen gas to synthesize propane.

Hydrogenation of Propene to Propane

Step 1

Here the double bond is reduced, and broken, along with the breakage of the H-H bond. The values for the bond enthalpies can be checked from the reference chart above.

Step 2

Sum of the bond enthalpies for C=C i.e. 610 kJ/mol and H-H bond i.e. 436 kJ/mol. The sum comes out to be 1046 kJ/mol.

Hydrogenation of Propene to Propane

Step 3

Here three new bonds are formed simultaneously. One bond between the C-C and two bonds between each hydrogen atom and carbon.

Step 4

The bond enthalpy of the C-C (-346 kJ/mol) and two C-H (-413 kJ/mol) bonds is determined and added together to determine the total energy released during the reaction. The sum is determined to be - 1172 kJ/mol.

Step 5

According to the formula of enthalpy of a reaction, ΔH_rxn, we get

ΔH_rxn  = energy added to break the bonds (+ ΔH) + energy released while forming the bonds (- ΔH)

⇒ ΔH_rxn  = (1046-1172) kJ/mol

⇒ ΔH_rxn  = -126 kJ/mol

Hence, the reduction of propene is an exothermic reaction.

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Things to Remember

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• Bond enthalpy is the amount of energy contained in a bond formed between atoms in a molecule.
• It is also known as bond-dissociation enthalpy, average bond energy, or bond strength.
• The bond enthalpy is the total amount of energy necessary to break one mole of that bond.
• Bond enthalpy gives an estimate and insight into the bond strength and its stability.
• The higher the bond enthalpy, the higher the energy required to break the bond, and the higher the bond strength.
• The value of bond enthalpy also determines the nature of the chemical process.

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