In general, compounds with negative \(\Delta H_{i}^{\circ}\) values are more stable than those with positive \(\Delta H_{i}^{\circ}\) values. \(\mathrm{H}_{2} \mathrm{O}_{2}(l)\) has a negative \(\Delta H_{\mathrm{f}}^{\circ}\) (see Table 6.4 ). Why, then, does \(\mathrm{H}_{2} \mathrm{O}_{2}(l)\) have a tendency to decompose to \(\mathrm{H}_{2} \mathrm{O}(l)\) and \(\mathrm{O}_{2}(g) ?\)

Enthalpy change, denoted as e , represents the heat exchange in a chemical reaction at constant pressure. It's a key concept in understanding chemical stability, as it indicates whether a compound will release or absorb heat during formation. A negative e implies that the reaction is exothermic, meaning that energy is released into the surroundings as the reactants transform into products. This release of energy tends to make the substance more stable, because systems generally favor a state of lower energy. Conversely, a positive e suggests an endothermic reaction, absorbing energy and resulting in a less stable compound.
However, enthalpy is just one aspect of stability. While e can predict whether a compound might prefer to exist, it does not give the complete picture, as there are other factors involved such as entropy and Gibbs free energy.

Gibbs free energy, often symbolized as G, is the ultimate arbiter of spontaneity and stability in chemical reactions. The change in Gibbs free energy, abla G , combines enthalpy, entropy, and temperature to predict whether a process will occur spontaneously. The formula abla G = abla H - T abla S where T is the temperature in Kelvin, can determine if a reaction is thermodynamically favorable.
abla G must be negative for a process to be spontaneous. Even if a reaction is exothermic (negative abla H ), it will not occur if the entropy decrease is sufficient to make abla G positive once multiplied by the temperature. Thus, the Gibbs free energy provides a more comprehensive view of reaction stability than enthalpy alone.

Entropy, symbolized by S, is often considered as a measure of disorder or randomness within a system. It is an essential part of understanding chemical reactions and stability. An increase in entropy suggests a move towards a more disordered state, which is generally favored in nature, as the universe tends towards maximum entropy. When a reaction leads to greater disorder, for example by producing more gas molecules or creating more possible arrangements for particles, the reaction is often spontaneous.
A decrease in entropy, indicating a shift towards order, can be a driving force for instability. Despite having a negative enthalpy change, a reaction could still be non-spontaneous if the entropy decrease is significant enough. Therefore, when assessing chemical stability, changes in entropy must be considered alongside changes in enthalpy to determine the overall free energy change of a system.

A redox reaction is a chemical process involving the transfer of electrons between two species. It's a combination of two simultaneous processes – oxidation, where a substance loses electrons, and reduction, where a substance gains electrons. These reactions are key to determining the stability of compounds, particularly in understanding why some substances have a tendency to decompose.
In the specific case of hydrogen peroxide (H₂O₂), the compound acts as both an oxidizer and a reducer, ultimately breaking down into water (H₂O) and oxygen (O₂). Even with a negative enthalpy of formation, the decomposition of H₂O₂ is favorable due to the significant increase in entropy associated with the production of gaseous oxygen and the stability of water as a product. This example clearly illustrates how redox reactions can contribute to the stability or instability of chemical compounds.