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Terminology: system, surroundings, thermal parameters State variables: average values of particular observables, such as pressure, elctric polarization, magnetization, displacement, the number of particles. Environmental parameters: volume, area, electric field E, magnetic induction field B, fixed particle number, temperature, che- mical potential μ. Two classes of Thermal parameters:

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Temperature: intuitively clear, but more difficult to define in precise physical terms (such a definition will emerge Later in the course). Equilibrium state: EXTREMELY important notion in Thermodynamics – in it’s classical as well as in quantum version! External interactions:

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Heat: energy transferred to the system through diathermal contact (a.k.a. diathermic contact). Very important: heat should not be confused with internal energy (conventionally, denoted as U ). The “Zeroth Law” of thermodynamics: Here, we already have to introduce the notion of tempera- ture. But we don’t yet know the definition! So, we will use a provisional “operational definition”: namely, T is what is measured by a thermometer.

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The “Zeroth Law”, cont.: Two or more systems in thermal contact and in thermal equilibrium all have the same temperature; If system A is in thermal equilibrium with system B, and system B is in thermal equilibrium with system C, then A and C are also in thermal equilibrium,

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The First Law: Where: Q – energy diathermally transferred to the system; and W – work done by the system. Be careful! This convention changes from book to book! In some textbooks the First Law is written as ΔU = Q + W. The First Law is extremely well summarized at the bottom of Page 8 in Dr, Wasserman’s text: The first Law simply states that a change in the internal energy of a system appears as the numerical difference between thermally dissipated energy and external electromechanical energy (work).

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The work and the heat depend on how the thermodynamic process is carried out. In contrast, the internal energy U is a measurable state variable. The differential dU is a true mathematical differential. The change ΔU in the system internal energy Depends only on the endpoints. The endpoints A and B represent macroscopic states of the system. An illustrative example:

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In our everyday lives we use many types of heat engines and devices that work in cycles. In a heat engine cycle, some heat goes in: Q in. Then Some heat goes out (Q out ), and some work is produced (W out ). Ifand were true differentials, the only way of returning to the initial state in the cycle would be if in = out, so no output work would be produced! There would be no heat engines (e.g., car engines), no refrigerators, no air conditioners!

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