Endothermic and Exothermic Processes

Every chemical reaction involves two energy steps: breaking bonds in the reactants and forming bonds in the products. Breaking a bond always costs energy — you have to pull atoms apart against their attraction, so energy flows in. Forming a bond always pays energy back — atoms settle into a lower-energy arrangement, so energy flows out. A reaction's overall heat change, delta H, is just the difference: energy absorbed to break old bonds minus energy released forming new bonds. If forming the new bonds releases more than breaking the old ones cost, the leftover energy escapes as heat, the surroundings warm up, and delta H is negative. That's exothermic. If breaking the old bonds costs more than the new bonds pay back, the reaction has to pull extra energy from the surroundings, the surroundings cool, and delta H is positive. That's endothermic.

Take methane combustion, CH₄ + 2O₂ → CO₂ + 2H_2O. Breaking the C-H and O=O bonds in the reactants takes energy, but forming the C=O and O-H bonds in the products releases far more. The surplus leaves as heat, which is why combustion feels hot and delta H is strongly negative. An instant cold pack works the opposite way: breaking apart the ionic lattice of a salt like ammonium nitrate as it dissolves costs more energy than the new ion-water interactions give back, so the pack pulls heat from your hand.

Reversing a reaction swaps which direction the energy flows, so it flips the sign of delta H exactly: if forming CO₂ and water releases 802 kJ/mol, decomposing them back into methane and oxygen requires absorbing that same 802 kJ/mol. This bond-energy bookkeeping sets up everything else in thermochemistry — energy diagrams, calorimetry, and Hess's law all build on tracking this same absorbed-versus-released balance.