Exothermic and endothermic reactions....
2H2(g) + O2(g) --> 2H2O(g) ........... exothermic..... heat is given off
N2(g) + O2(g) --> 2NO(g) ............... endothermic... heat is absorbed
The majority of chemical reactions are exothermic, that is, they have negative changes in enthalpy. ΔH is negative. Endothermic reactions have positive enthalpy changes. ΔH is positive.
Contrary to what some teachers might say, reactions don't occur so that the products will have "stable octets."
Instead, chemical reactions are all about energy. For example, place a bowling ball on a smooth hill. How does the ball move? Of course the ball rolls downhill. That's because spontaneous processes (those with no outside intervention) result in a lower energy. In the case of the bowling ball, it moves to a lower gravitational potential energy.
Chemical reactions move to a lower overall energy. Often that means that the thermal energy of the products is less than the thermal energy of the reactants. This translates to a negative change in enthalpy. These are exothermic reactions, and they are easy to compare to bowling balls on the sides of hills.
But what about an endothermic reaction where the product is at a higher thermal energy, that is, a greater enthalpy? Bowling balls don't roll uphill all on their own.
In chemistry the overall energy is the combination of enthalpy, ΔH, and entropy, ΔS. Entropy is a measure of the number of microstates in which substances exist. The connection between enthalpy and entropy is given by the equation for Gibbs free energy, ΔG.
ΔG = ΔH - TΔS
When ΔG is negative the reaction is spontaneous, that is, the bowling ball rolls downhill. The reaction moves to a lower total energy. That is possible even if ΔH is positive (endothermic) because of the increase in entropy. When ΔS is positive and ΔH is positive, ΔG can still be negative, and the reaction be spontaneous. The positive ΔH is offset by the -TΔS term.