fayled said:
I still don't feel comfortable with the way I have calculated the reservoir entropy change though. For 1. the reservoir entropy change is not reversible yet I have used my formula which was derived from using the definition of entropy dS=∫dQrev/T.
You calculated the entropy change correctly for the reservoirs in all the cases. You indicate that, for problem 1, "the reservoir entropy change is not reversible." I think you mean is that "the heat transfer process for the reservoir was not reversible." Actually, you don't know this.
You have enough knowledge right now to correctly calculate the
change in entropy for the combined system of water plus reservoir(s), but you currently don't have enough knowledge to ascertain in which part of the combined system the main amount of irreversibility occurred. That all depends on the transient temperature behavior of the system. And in particular, it depends on the temperature vs time variation and the rate of heat flow variation at the interface between the water and the reservoir(s). When we say that a reservoir is a constant temperature reservoir, we mean that,
in the initial and final equilibrium states of the combined system, the reservoir temperature is the same (and uniform). But, during the time that heat is being transferred between the water and the reservoir, the temperature within the reservoir will not be uniform, and will be hotter near the interface with the hotter water. And, during the time that the heat is being transferred, the temperature of the water will not be uniform, but will be colder near the interface than the colder reservoir. This is how heat transfer by heat conduction, an irreversible exchange, occurs (remember from freshman physics).
Have you learned yet about the Clausius inequality. Clausius determined that, for an irreversible process applied to a closed system that
\Delta S > \int{\frac{dQ}{T_I}}
where T
I is the temperature at the interface (boundrary) of the system and dQ is the heat transferred at the boundary. This inequality applies to the water in our system as well as to the reservoir(s). Each of these are a subsystem of the combined system. If you want to get a handle on the amount of irreversibility that occurred in a particular subsystem, you just take \Delta S and subtract \int{\frac{dQ}{T_I}}. The larger this number, the more the amount of irreversibility that took place in that subsystem.
If we evaluate the above inequality for both the water and the reservoirs, we get:
\Delta S_w > \left(\int{\frac{dQ}{T_I}}\right)_w
\Delta S_r > \left(\int{\frac{dQ}{T_I}}\right)_r
But we know that the heat flow to the reservoir is equal to minus the heat flow to the water, and the temperature at the interface between the water and the reservoir is the same for both. So, if we add the two inequalities together, we get:
\Delta S_w + \Delta S_r > 0
This is what we found, irrespective of the transient heat transfer behavior of the combined system.
I would also like note that, from your previous calculations, as you add more reservoirs to the sequence, the combined entropy change decreases inversely with the number of reservoirs (as you can show by further calculations).
Also, here is a short write up on the first and second laws that may be helpful for your understanding:FIRST LAW OF THERMODYNAMICS
Suppose that we have a closed system that at initial time t
i is in an initial equilibrium state, with internal energy U
i, and at a later time t
f, it is in a new equilibrium state with internal energy U
f. The transition from the initial equilibrium state to the final equilibrium state is brought about by imposing a time-dependent heat flow across the interface between the system and the surroundings, and a time-dependent rate of doing work at the interface between the system and the surroundings. Let \dot{q}(t) represent the rate of heat addition across the interface between the system and the surroundings at time t, and let \dot{w}(t) represent the rate at which the system does work on the surroundings at the interface at time t. According to the first law (basically conservation of energy),
\Delta U=U_f-U_i=\int_{t_i}^{t_f}{(\dot{q}(t)-\dot{w}(t))dt}=Q-W
where Q is the total amount of heat added and W is the total amount of work done by the system on the surroundings at the interface.
The time variation of \dot{q}(t) and \dot{w}(t) between the initial and final states uniquely characterizes the so-called process path. There are an infinite number of possible process paths that can take the system from the initial to the final equilibrium state. The only constraint is that Q-W must be the same for all of them.
If a process path is irreversible, then the temperature and pressure within the system are inhomogeneous (i.e., non-uniform, varying with spatial position), and one cannot define a unique pressure or temperature for the system (except at the initial and the final equilibrium state). However, the pressure and temperature
at the interface can be measured and controlled using the surroundings to impose the temperature and pressure boundary conditions that we desire. Thus, T
I(t) and P
I(t) can be used to impose the process path that we desire. Alternately, and even more fundamentally, we can directly control, by well established methods, the rate of heat flow and the rate of doing work at the interface \dot{q}(t) and \dot{w}(t)).
Both for reversible and irreversible process paths, the rate at which the system does work on the surroundings is given by:
\dot{w}(t)=P_I(t)\dot{V}(t)
where \dot{V}(t) is the rate of change of system volume at time t. However, if the process path is reversible, the pressure P within the system is uniform, and
P_I(t)=P(t) (reversible process path)
Therefore, \dot{w}(t)=P(t)\dot{V}(t) (reversible process path)
Another feature of reversible process paths is that they are carried out very slowly, so that \dot{q}(t) and \dot{w}(t) are both very close to zero over then entire process path. However, the amount of time between the initial equilibrium state and the final equilibrium state (t
f-t
i) becomes exceedingly large. In this way, Q-W remains constant and finite.
SECOND LAW OF THERMODYNAMICS
In the previous section, we focused on the infinite number of process paths that are capable of taking a closed thermodynamic system from an initial equilibrium state to a final equilibrium state. Each of these process paths is uniquely determined by specifying the heat transfer rate \dot{q}(t) and the rate of doing work \dot{w}(t) as functions of time at the interface between the system and the surroundings. We noted that the cumulative amount of heat transfer and the cumulative amount of work done over an entire process path are given by the two integrals:
Q=\int_{t_i}^{t_f}{\dot{q}(t)dt}
W=\int_{t_i}^{t_f}{\dot{w}(t)dt}
In the present section, we will be introducing a third integral of this type (involving the heat transfer rate \dot{q}(t)) to provide a basis for establishing a precise mathematical statement of the Second Law of Thermodynamics.
The discovery of the Second Law came about in the 19th century, and involved contributions by many brilliant scientists. There have been many statements of the Second Law over the years, couched in complicated language and multi-word sentences, typically involving heat reservoirs, Carnot engines, and the like. These statements have been a source of unending confusion for students of thermodynamics for over a hundred years. What has been sorely needed is a precise mathematical definition of the Second Law that avoids all the complicated rhetoric. The sad part about all this is that such a precise definition has existed all along. The definition was formulated by Clausius back in the 1800's.
Clausius wondered what would happen if he evaluated the following integral over each of the possible process paths between the initial and final equilibrium states of a closed system:
I=\int_{t_i}^{t_f}{\frac{\dot{q}(t)}{T_I(t)}dt}
where T
I(t) is the temperature at the interface with the surroundings at time t. He carried out extensive calculations on many systems undergoing a variety of both reversible and irreversible paths and discovered something astonishing. He found that, for any closed system, the values calculated for the integral over all the possible reversible and irreversible paths (between the initial and final equilibrium states) was not arbitrary; instead, there was a unique upper bound (maximum) to the value of the integral. Clausius also found that this result was consistent with all the "word definitions" of the Second Law.
Clearly, if there was an upper bound for this integral, this upper bound had to depend only on the two equilibrium states, and not on the path between them. It must therefore be regarded as a point function of state. Clausius named this point function Entropy.
But how could the value of this point function be determined without evaluating the integral over every possible process path between the initial and final equilibrium states to find the maximum? Clausius made another discovery. He determined that, out of the infinite number of possible process paths, there existed a well-defined subset, each member of which gave the same maximum value for the integral. This subset consisted of what we call today
the reversible process paths. So, to determine the
change in entropy between two equilibrium states, one must first conceive of a reversible path between the states and then evaluate the integral. Any other process path will give a value for the integral lower than the entropy change.
So, mathematically, we can now state the Second Law as follows:
I=\int_{t_i}^{t_f}{\frac{\dot{q}(t)}{T_I(t)}dt}\leq\Delta S=\int_{t_i}^{t_f} {\frac{\dot{q}_{rev}(t)}{T(t)}dt}
where \dot{q}_{rev}(t) is the heat transfer rate for any of the reversible paths between the initial and final equilibrium states, and T(t) is the
system temperature at time t (which, for a reversible path, is equal to the temperature at the interface with the surroundings). This constitutes a precise mathematical statement of the Second Law of Thermodynamics.
Chet
