In chemical reactions, understanding the bonds that are broken and formed is crucial for determining the energy changes involved. When analyzing a reaction, the bonds on the reactants side represent those that are broken, while the bonds on the products side represent those that are formed. For instance, if a carbon-hydrogen (C-H) bond is broken and replaced with a carbon-bromine (C-Br) bond, it indicates that energy is required to break the initial bond, leading to a positive enthalpy change for that process.

Conversely, the formation of new bonds releases energy, resulting in a negative enthalpy change. This distinction is essential for calculating the overall enthalpy change (ΔH) of the reaction. The enthalpy change can be calculated using the bond dissociation energies of the bonds involved. For example, if the C-H bond has a bond energy of 436 kJ/mol and the bromine-bromine (Br-Br) bond has a bond energy of 192 kJ/mol, these values would be added positively since they represent energy input for breaking bonds.

On the products side, if the newly formed C-Br bond has a bond energy of 923 kJ/mol and the hydrogen-bromine (H-Br) bond has a bond energy of 368 kJ/mol, these values would be subtracted as they represent energy released during bond formation. The overall enthalpy change can be calculated as follows:

ΔH = (Bond energies of broken bonds) - (Bond energies of formed bonds)

Substituting the values gives:

ΔH = (436 + 192) - (923 + 368) = -33 kJ/mol

A negative ΔH indicates that the reaction is exothermic, meaning it releases energy and results in a more stable product state. However, to determine if the reaction is spontaneous, additional information such as temperature and the change in entropy (ΔS) is required. Thus, while we can conclude that the reaction is exothermic, further analysis is needed to assess spontaneity.

Struggling with Organic Chemistry?