Introduction to entropy: Difference between revisions

Revision as of 20:46, 8 July 2024

Revision as of 20:47, 8 July 2024

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Ice melting provides an example in which entropy increases in a small system, a thermodynamic system consisting of the surroundings (the warm room) and the entity of glass container, ice and water which has been allowed to reach [[thermodynamic equilibrium]] at the melting temperature of ice. In this system, some [[heat]] (''δQ'') from the warmer surroundings at 298 K (25 °C; 77 °F) transfers to the cooler system of ice and water at its constant temperature (''T'') of 273 K (0 °C; 32 °F), the melting temperature of ice. The entropy of the system, which is {{sfrac|δ''Q''|''T''}}, increases by {{sfrac|δ''Q''|273 K}}. The heat δ''Q'' for this process is the energy required to change water from the solid state to the liquid state, and is called the [[enthalpy of fusion]], i.e. Δ''H'' for ice fusion.

It is important to realize that theThe entropy of the surrounding room decreases less than the entropy of the ice and water increases: the room temperature of 298 K is larger than 273 K and therefore the ratio, (entropy change), of {{sfrac|δ''Q''|298 K}} for the surroundings is smaller than the ratio (entropy change), of {{sfrac|δ''Q''|273 K}} for the ice and water system. This is always true in spontaneous events in a thermodynamic system and it shows the predictive importance of entropy: the final net entropy after such an event is always greater than was the initial entropy.

As the temperature of the cool water rises to that of the room and the room further cools imperceptibly, the sum of the {{sfrac|δ''Q''|''T''}} over the continuous range, "at many increments", in the initially cool to finally warm water can be found by calculus. The entire miniature 'universe', i.e. this thermodynamic system, has increased in entropy. Energy has spontaneously become more dispersed and spread out in that 'universe' than when the glass of ice and water was introduced and became a 'system' within it.