At 0 degrees Celsius, the ion product of water, denoted as \( K_w \), is recorded as \( 1.2 \times 10^{-15} \). This value indicates the concentration of hydronium ions (\( H_3O^+ \)) and hydroxide ions (\( OH^- \)) in a neutral solution at this temperature. As temperature increases to 25 degrees Celsius, \( K_w \) rises to \( 1.0 \times 10^{-14} \). This temperature dependence of \( K_w \) is crucial in understanding the behavior of water and its autoionization process, represented by the equation:

\[ H_2O_{(l)} + H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + OH^-_{(aq)} \]

According to Le Chatelier's principle, an increase in temperature shifts the equilibrium position away from heat. In this case, since \( K_w \) increases with temperature, it suggests that the formation of products (\( H_3O^+ \) and \( OH^- \)) is favored. This shift indicates that the reaction is endothermic, meaning heat is absorbed as a reactant in the process. Therefore, the reaction can be viewed as:

\[ \text{Heat} + H_2O_{(l)} + H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + OH^-_{(aq)} \]

In contrast, if the temperature were to decrease, the equilibrium would shift to the left, favoring the reactants and resulting in a decrease in \( K_w \). Under thermal neutral conditions, where the change in enthalpy (\( \Delta H \)) is zero, \( K_w \) remains constant, indicating no shift in equilibrium. Thus, understanding the relationship between temperature and \( K_w \) not only helps in determining the direction of the reaction but also in classifying the reaction as endothermic or exothermic.

In summary, the increase in \( K_w \) with temperature signifies that the reaction is endothermic, aligning with the principles of chemical equilibrium and thermodynamics.