Introduction to entropy: Difference between revisions

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== Example of increasing entropy ==

{{Main article|Disgregation}}

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 {{Convert|298 |K (25 °|C; 77 °F)}} transfers to the cooler system of ice and water at its constant temperature (''T'') of {{Convert|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.

The 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.