## 3.4 Bond Energy (Bond Enthalpy) **Bond energy**, also known as *bond enthalpy*, is the energy required to break one mole of a specific type of bond in a gaseous substance. It is a measure of the strength of a chemical bond. - **Unit**: Kilojoules per mole ($kJ\,mol^{-1}$) Bond energy is influenced by several factors: - **Electronegativity difference** between the two bonded atoms - **Size of the atoms** - **Bond order** (single, double, triple bond) - **Bond length** ### Effect of Electronegativity The greater the electronegativity difference between two bonded atoms, the more polar the bond becomes, which in turn increases the bond energy. A more polar bond has stronger electrostatic attraction, requiring more energy to break. This trend is clearly visible in the hydrogen halides (hydrides of Group VII). **Table 3.2: Bond Energy of Hydrogen Halides** | Bond | Bond Energy ($kJ\,mol^{-1}$) | Electronegativity Difference | | :--- | :--- | :--- | | $H-F$ | 562 | 1.8 | | $H-Cl$| 431 | 0.9 | | $H-Br$| 366 | 0.7 | | $H-I$ | 299 | 0.4 | From the data, as the electronegativity difference decreases from $H-F$ to $H-I$, the bond energy also decreases. $H-F$ has the largest electronegativity difference and the highest bond energy. ### Bond Length **Bond length** is the average distance between the nuclei of two covalently bonded atoms. It is typically measured in Ångströms ($Å$) or picometers ($pm$). $$1 \, pm = 10^{-12} \, m$$ #### Relationship with Atomic Size and Bond Energy Bond length is directly related to the size of the bonded atoms. Generally, smaller atoms form shorter bonds, and shorter bonds are stronger, meaning they have higher bond energy. **Example**: The bond energy of $H-H$ is $436 \, kJ\,mol^{-1}$, while that of $H-Br$ is $366 \, kJ\,mol^{-1}$. This is because the hydrogen molecule ($H_2$) has a much shorter bond length compared to hydrogen bromide ($HBr$), as the bromine atom is significantly larger than a hydrogen atom. #### Bond Properties of Halogens The relationship between bond length and bond energy can be observed in the halogen group. For more details on these elements, see . **Table 3.3: Bond Energy and Bond Length of Halogens** | Bond | Bond Energy ($kJ\,mol^{-1}$) | Bond Length (pm) | | :--- | :--- | :--- | | $F-F$ | 158 | 149 | | $Cl-Cl$| 242 | 199 | | $Br-Br$| 193 | 228 | | $I-I$ | 151 | 266 | As the atomic radius increases down the group from fluorine to iodine, the bond length increases. The bond energy generally decreases from chlorine to iodine. **The Fluorine Anomaly**: An exception to the trend is fluorine ($F_2$), which has a lower bond energy than chlorine ($Cl_2$). This is because fluorine atoms are very small and electron-dense. The lone pairs of electrons on the adjacent fluorine atoms repel each other strongly, which weakens the $F-F$ covalent bond. Consequently, less energy is required to break it. ### Predicting Reactivity Bond energy is a useful tool for predicting the reactivity of covalent molecules. - **High Bond Energy**: A high bond energy indicates a strong bond. Molecules with strong bonds are more stable and therefore less reactive. - **Low Bond Energy**: A low bond energy indicates a weak bond. Molecules with weak bonds are less stable and more reactive. **Example**: In alkyl halides, the $C-F$ bond has a greater bond energy than the $C-Cl$ bond. Therefore, alkyl fluorides are generally less reactive than alkyl chlorides.