Explain why exothermic processes are mire likely spontaneous when temperature of the system is low?
Exothermic processes are more likely to be spontaneous at low temperatures because of the relationship between enthalpy change and entropy change.
To understand this, we need to consider the two main factors that determine spontaneity: enthalpy (ΔH) and entropy (ΔS).
Enthalpy (ΔH) refers to the heat energy exchanged during a process. In exothermic processes, the reaction releases heat energy, resulting in a negative enthalpy change (ΔH < 0).
Entropy (ΔS), on the other hand, refers to the level of disorder or randomness in a system. Processes that increase the level of disorder have a positive entropy change (ΔS > 0).
The concept of spontaneity is governed by the Gibbs Free Energy equation: ΔG = ΔH - TΔS, where ΔG is the change in Gibbs free energy and T is the temperature.
At low temperatures, the TΔS term in the equation becomes relatively small compared to the enthalpy term. Since exothermic processes have a negative enthalpy change (ΔH < 0), the negative enthalpy term dominates the equation, making the overall ΔG negative.
A negative ΔG indicates that the reaction is spontaneous under these conditions. In other words, the exothermic process is more likely to occur spontaneously, releasing energy to the surroundings.
However, as the temperature increases, the TΔS term becomes larger and can potentially counterbalance the negative enthalpy term. At higher temperatures, the positive entropy change (ΔS > 0) can overcome the negative enthalpy change (ΔH < 0), resulting in a positive ΔG.
Therefore, at higher temperatures, exothermic processes are less likely to be spontaneous because the increase in disorder becomes significant enough to outweigh the release of heat energy.
It is important to mention that the explanation above provides a general understanding, but specific cases may have additional factors influencing spontaneity.