What is 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.
FAQs on Bond Energy
1. What is bond energy in chemistry?
Bond energy, also known as bond enthalpy, is the average amount of energy required to break one mole of a specific type of bond between two atoms in the gaseous state. It serves as a key measure of the strength and stability of a chemical bond. A higher bond energy value indicates a stronger and more stable bond that requires more energy to break.
2. What is the main difference between bond energy and bond dissociation energy?
The key difference lies in their specificity. Bond dissociation energy is the exact energy needed to break one specific bond in a particular molecule. In contrast, bond energy is the average value calculated from the bond dissociation energies of all bonds of the same type within a molecule. For example, in methane (CH₄), the energy to break the first C-H bond is different from the second, but the C-H bond energy is the average of all four.
3. How is the average bond energy calculated for a polyatomic molecule?
To calculate the average bond energy for a specific type of bond in a polyatomic molecule, you must first determine the total energy required to break all bonds of that type to separate the atoms. This total energy is then divided by the number of bonds broken. For a molecule like water (H₂O), the average O-H bond energy is the sum of the energies needed to break the first and second O-H bonds, divided by two.
4. What are the key factors that influence the strength of a chemical bond?
Several factors determine the bond energy, or strength, of a chemical bond:
- Atomic Size: Smaller atoms form shorter and stronger bonds because their nuclei can get closer, leading to greater electrostatic attraction.
- Bond Order (Multiplicity): The number of shared electron pairs affects strength. A triple bond (e.g., in N₂) is stronger than a double bond (e.g., in O₂), which is stronger than a single bond (e.g., in F₂).
- Electronegativity Difference: A larger difference in electronegativity between two atoms results in a more polar bond, which generally increases the bond strength due to added ionic character.
- Lone Pair Repulsion: The presence of lone pairs of electrons on adjacent bonded atoms can cause repulsion, which weakens the bond.
5. How does bond order affect bond energy? Explain with an example.
Bond order has a direct and significant impact on bond energy: as the bond order increases, the bond energy also increases. This is because more electron pairs are shared between the two atoms, holding them together more tightly. For example, the nitrogen molecule (N₂) has a triple bond (bond order = 3) and a very high bond energy of 945 kJ/mol. In comparison, the oxygen molecule (O₂) has a double bond (bond order = 2) with a lower bond energy of 498 kJ/mol.
6. Why is bond energy considered an average value and not an exact one?
Bond energy is an average because the energy required to break a particular type of bond is not constant and depends on its molecular environment. For instance, in a water molecule (H-O-H), breaking the first H-O bond to form H and OH radicals requires a different amount of energy than breaking the O-H bond in the resulting OH radical. Since the chemical environment changes after each bond is broken, we use the average of these successive bond dissociation energies to represent a typical bond strength.
7. How can atomic radius be used to predict the relative strength of a bond?
There is an inverse relationship between atomic radius and bond strength. Generally, smaller atomic radii lead to shorter bond lengths. A shorter bond allows for greater overlap between electron orbitals, resulting in a stronger electrostatic attraction between the nuclei and the shared electrons. This increased attraction leads to a higher bond energy. For instance, the H-F bond is significantly stronger than the H-I bond because fluorine has a much smaller atomic radius than iodine.
8. Which typically has a higher bond energy: an ionic bond or a covalent bond, and why?
Generally, ionic bonds have higher bond energies than covalent bonds. This is because ionic bonds are formed from the strong, non-directional electrostatic attraction between oppositely charged ions that are arranged in a rigid crystal lattice. This lattice structure maximises attraction and minimises repulsion, requiring a large amount of energy (lattice energy) to break apart. While some covalent bonds like N≡N are very strong, most covalent compounds, which involve electron sharing, have lower melting and boiling points, reflecting weaker overall forces and lower bond energies compared to ionic compounds.
