Chemistry – Chemical energetics - Exothermic and endothermic reactions | e-Consult

The student's statement is incorrect. While it is true that bond breaking is generally endothermic and bond making is generally exothermic, it is not accurate to say that all reactions are exothermic. Here's a critical evaluation:

Why the statement is incorrect:

• Enthalpy change depends on the relative energies of bonds broken and formed. The enthalpy change of a reaction is determined by the difference in energy between the bonds broken and the bonds formed, not simply whether bond breaking is endothermic and bond making is exothermic.
• Some reactions require a significant amount of energy input for bond breaking, even if the bond formation releases a large amount of energy. If the energy required to break the initial bonds is greater than the energy released when new bonds form, the overall reaction will be endothermic.

Examples to support the counter-argument (endothermic reactions):

• Photosynthesis: CO₂(g) + H₂O(l) → C₆H₁₂O₆(aq) + O₂(g) is an endothermic reaction. While the formation of bonds in glucose and oxygen releases energy, the energy required to break the bonds in CO₂ and H₂O is greater. Photosynthesis requires a constant input of light energy to proceed.
• Decomposition of Calcium Carbonate (CaCO₃): CaCO₃(s) → CaO(s) + CO₂(g) is an endothermic reaction. Breaking the strong covalent bonds in CaCO₃ requires a significant amount of energy, and the energy released when forming CaO and CO₂ is not sufficient to overcome this. This is why calcium carbonate decomposes when heated.
• Dissolving some salts: The dissolution of certain ionic compounds in water can be endothermic, requiring energy input to break the ionic lattice and hydrate the ions.

In conclusion, while bond breaking is typically endothermic and bond making is typically exothermic, the overall enthalpy change of a reaction depends on the balance between the energies of the bonds broken and the bonds formed. Therefore, not all reactions are exothermic.