Predict the signs of \(\Delta H, \Delta S\), and \(\Delta G\) of the system for the following processes at 1 atm: (a) ammonia melts at \(-60^{\circ} \mathrm{C},\) (b) ammonia melts at \(-77.7^{\circ} \mathrm{C}\), (c) ammonia melts at \(-100^{\circ} \mathrm{C}\). (The normal melting point of ammonia is \(-77.7^{\circ} \mathrm{C} .\) )

When a substance like ammonia transitions from a solid to a liquid, its particles begin to move more freely, as they're no longer held in a rigid structure. This increase in randomness and disorder within the system is what we describe as an increase in entropy, denoted by a positive change in entropy, (∆S). In melting, the molecules are less confined and can explore more states than in their solid form; hence the entropy invariably increases. This concept explains why (∆S) is positive in all instances of ammonia melting, whether below, at, or above its melting point.

• Lower temperatures mean ammonia molecules move less energetically, but still more than in solid form.
• At or above the melting point, the increase in molecular movement—and thus entropy—is even more pronounced.

Entropy is a fundamental concept in understanding not just phase changes but the direction of natural processes as a whole, as it provides a measure of the energy dispersal in a given system.

Phase transitions, such as melting, require energy changes within the substance. In the case of ammonia, as it melts, it absorbs heat from its surroundings, an endothermic process that results in an increase in heat content, or enthalpy, of the system. This is universally true for melting processes and is represented by a positive change in enthalpy, (∆H). These heat exchanges are pivotal in categorizing phase transitions and understanding energy flows within physical systems.

• Melting requires an input of energy to overcome intermolecular forces.
• The endothermic nature of melting results in a positive (∆H).

Exploring these energy changes offers insights into the energetic requirements for phase transitions and the inherent properties of substances under different conditions.

The Gibbs Free Energy, (∆G), helps predict the spontaneity of a phase transition. It illustrates whether a process will occur without external energy input. A negative (∆G) denotes a spontaneous process, positive for non-spontaneous, and zero when the system is in equilibrium. For ammonia:

• Below its melting point (-77.7°C), the process of melting is non-spontaneous ((∆G) is positive).
• At its melting point, ammonia is in equilibrium and (∆G) is zero, indicating no net drive for melting or freezing.
• Above (-77.7°C), melting is spontaneous, as reflected by a negative (∆G).

This criterion of spontaneous change is crucial for understanding the natural direction of many physical and chemical processes and the inherent chemical potential of substances.

The melting point of a substance is the precise temperature at which it transitions from solid to liquid at atmospheric pressure. For ammonia, this temperature is (-77.7°C) under standard atmospheric conditions (1 atm). The melting point is intrinsic to the substance, representing the balance between kinetic energy and intermolecular forces at which the solid and liquid phases coexist in equilibrium.

• At temperatures below (-77.7°C), solid ammonia requires additional thermal energy to melt.
• At the melting point, ammonia exists as both a solid and a liquid, reflecting a state of dynamic equilibrium.
• Above this temperature, ammonia will naturally progress towards a liquid state as it possesses sufficient kinetic energy to overcome the forces maintaining its solid structure.

Understanding the melting point is essential in predicting behavior and designing systems involving ammonia, whether in industrial applications or scientific research.