Lesson Plan

Describe entropy as a measure of disorder, of energy unavailable to do work, and of the dispersal of energy

Calculate the change in entropy using ΔS = Q/T, with correct signs and absolute (Kelvin) temperature

State the second law of thermodynamics and explain why heat transfers spontaneously from hot to cold but never spontaneously from cold to hot

Explain how the entropy of a local system can decrease without violating the second law (ΔS_tot = ΔS_syst + ΔS_envir > 0)

12.3

"Entropy is a measure of the disorder of a system. Entropy also describes how much energy is not available to do work. The more disordered a system and higher the entropy, the less of a system's energy is available to do work." "The equation for the change in entropy, $\Delta S$, is $\Delta S = \frac{Q}{T}$, where Q is the heat that transfers energy during a process, and T is the absolute temperature at which the process takes place." "The second law of thermodynamics states that the total entropy of a system either increases or remains constant in any spontaneous process; it never decreases. An important implication of this law is that heat transfers energy spontaneously from higher- to lower-temperature objects, but never spontaneously in the reverse direction." "Based on this equation, we see that $\Delta S_{\text{syst}}$ can be negative as long as $\Delta S_{\text{envir}}$ is positive and greater in magnitude." (referring to $\Delta S_{\text{tot}} = \Delta S_{\text{syst}} + \Delta S_{\text{envir}} > 0$) "Find the increase in entropy of 1.00 kg of ice that is originally at 0 °C and melts to form water at 0 °C. … $\Delta S = \frac{Q}{T} = \frac{3.34 \times 10^5\text{ J}}{273\text{ K}} = 1.22 \times 10^3\text{ J/K}.$"