Thermodynamics system, definition

Consider two distinct closed thermodynamic systems each consisting of n moles of a specific substance in a volnme Vand at a pressure p. These two distinct systems are separated by an idealized wall that may be either adiabatic (lieat-impemieable) or diathermic (lieat-condncting). Flowever, becanse the concept of heat has not yet been introdnced, the definitions of adiabatic and diathemiic need to be considered carefiilly. Both kinds of walls are impemieable to matter a permeable wall will be introdnced later. 

We have previously emphasized (Section 2.10) the importance of considering only intensive properties Rt (rather than size-dependent extensive properties Xt) as the proper state descriptors of a thermodynamic system. In the present discussion of heterogeneous systems, this issue reappears in terms of the size dependence (if any) of individual phases on the overall state description. As stated in the caveat regarding the definition (7.7c), the formal thermodynamic state of the heterogeneous system is wholly / dependent of the quantity or size of each phase (so long as at least some nonvanishing quantity of each phase is present), so that the formal state descriptors of the multiphase system again consist of intensive properties only. We wish to see why this is so. 

This section reviews some basic definitions and formulas in thermodynamics. These definitions will be used to develop energy balances to describe cooling tower operations. In our discussions we will use the following terms system, property, extensive and intensive properties, and... 

In summary, a reference state or standard state must be defined for each component in the system. The definition may be quite arbitrary and may be defined for convenience for any thermodynamic system, but the two states cannot be defined independently. When the reference state is defined, the standard state is determined conversely, when the standard state is defined, the reference state is determined. There are certain conventions that have been developed through experience but, for any particular problem, it is not necessary to hold to these conventions. These conventions are discussed in the following sections. The general practice is to define the reference state. This state is then a physically realizable state and is the one to which experimental measurements are referred. The standard state may or may not be physically realizable, and in some cases it is convenient to speak of the standard state for the chemical potential, for the enthalpy, for the entropy,... 

Throughout the discussions in Sections 8.15-8.18, we have emphasized methods for obtaining expressions for the chemical potential of a component when we choose to treat the thermodynamic systems in terms of the species that may be present in solution. A complete presentation of all possible types of systems containing charged or neutral molecular entities is not possible. However, no matter how complicated the system is, the pertinent equations can always be developed by the use of the methods developed here, together with the careful definition of reference states or standard states. We should also recall at this point that it is the quantity (nk — nf) that is determined directly or indirectly from experiment. 

The heat capacities that have been discussed previously refer to closed, single-phase systems. In such cases the variables that define the state of the system are either the temperature and pressure or the temperature and volume, and we are concerned with the heat capacities at constant pressure or constant volume. In this section and Section 9.3 we are concerned with a more general concept of heat capacity, particularly the molar heat capacity of a phase that is in equilibrium with other phases and the heat capacity of a thermodynamic system as a whole. Equation (2.5), C = dQ/dT, is the basic equation for the definition of the heat capacity which, when combined with Equation (9.1) or (9.2), gives the relations by which the more general heat capacities can be calculated. Actually dQ/dT is a ratio of differentials and has no value until a path is defined. The general problem becomes the determination of the variables to be used in each case and of the restrictions that must be placed on these variables so that only the temperature is independent. 

The gravitational field of the Earth is characterized by a potential, , that has a definite value at each point in the field. For all practical purposes this field is independent of the presence of matter in the quantities used in normal thermodynamic systems. Within this approximation the field is independent of the state of a thermodynamic system within it. The potential can be written as... 

By definition, the thermodynamic system defined in Figure 7.7 is an adiabatic system, then Ej = Ej. Therefore, combining Equations 7.12 and 7.13, and making the following approximations [40]... 

The Gibbs free energy is a measure of the probability that a reaction occurs. It is composed of the enthalpy, H, and the entropy, S° (Eq. 5). The enthalpy can be described as the thermodynamic potential, which ensues H = U + p V. where U is the internal energy, p is the pressure, and V is the volume. The entropy, according to classical definitions, is a measure of molecular order of a thermodynamic system and the irreversibility of a process, respectively. 

The reversibility of molecular behavior gives rise to a kind of symmetry in which the transport processes are coupled to each other. Although a thermodynamic system as a whole may not be in equilibrium, the local states may be in local thermodynamic equilibrium all intensive thermodynamic variables become functions of position and time. The definition of energy and entropy in nonequilibrium systems can be expressed in terms of energy and entropy densities u(T,Nk) and s(T,Nk), which are the functions of the temperature field T(x) and the mole number density Y(x) these densities can be measured. The total energy and entropy of the system is obtained by the following integrations... 

Give precise definitions for the terms thermodynamic system, open system, closed system, thermodynamic state, and reversible and irreversible process (Section 12.1). 

A thermodynamic system is in equilibrium when there is no change of intensive variables within the system. An equilibrium state is a state that cannot be changed without interactions with the environment. This definition includes that of mechanical equilibrium, but is a more general one. A state of thermodynamic equilibrium is a state of simultaneous chemical-, thermal-, and mechanical equilibrium. In other words, thermodynamic equihbrium is the state of the simultaneous vanishing of all fluxes ([32], p. 267). A thermodynamic system is thus... 

There are two hypothetical limiting cases of interest. In one, an infinitely slow cooling rate maintains thermodynamic equilibrium to the ideal glass, and the equilibrium formalism is applicable. In the other a fluid in equilibrium (at its fictive temperature) is quenched infinitely fast to a temperature low enough so that no molecular transport occurs. In this case, what were dynamic fluctuations in time becomes static fluctuations in space. The most elementary treatment of this glass is then as a thermodynamic system with one additional parameter, the fictive temperature. In an actual experiment, of course, relaxations take place and the state of the system is dependent upon its entire thermal history and requires many parameters for its definition. Detailed discussion of the use of irreversible thermodynamics for the study of relaxation processes in liquids and glasses is contained in reviews by Davies (1956, 1960). 

For convenience, thermodynamic systems are usually assumed closed, isolated from the surroundings. The laws that govern such systems are written in terms of two types of variables intensive (or intrinsic) that do not depend on the mass and extensive that do. By definition, extensive variables are additive, that is, their value for the whole system is the sum of their values for the individual parts. For example, volume, entropy, and total energy of a system are extensive variables, but the specific volume (or its reciprocity - the density), molar volume, or molar free energy of mixing are intensive. It is advisable to use, whenever possible, intensive variables. 

To begin with, Reiss (1965) has the conventional definition of a thermodynamic system (p. 3) ... 

The definition of a thermodynamic system in nanoscale is the same as the macroscopic systems. In thermod5mamics, a system is any region completely enclosed within a well-defined boundary. Everything outside the system is then defined as the surroundings. The boundary may be either... 

Thermodynamics which deals with systems that is essentially away from the equilibrium state is known as non-equilibrium thermodynamics. Many physical and chemical systems in the practical life are not in equilibrium state because they are continuously/discontinuously performing some actions where changes in their mass and energy from one form to another are taken place. Studies of such thermodynamical systems are required some added functionality other than basics of thermodynamic equilibrium systems. Several physical and chemical systems are remained beyond the reach of macroscopic methods of thermodynamics even nowadays. A major complexity for macroscopic thermodynamics is the definition of entropy which not comes exactiy in the thermodynamic equilibrium [3, 5]. 

The definition of the surface energy /emerges from a more general formulation of the Helmholtz or Gibbs free energy change of a thermodynamic system in the presence of surface effects ... 

The set of molecules of a fluid phase, i. e. a gas or a liquid, which is within the field of forces of the atoms or molecules of the external or internal surface of a solid sorbent is called the adsorbate of the fluid on this sorbent material. This set may be considered as a thermodynamic system or phase in the sense of W. Schottky [1.63]. Such a system is defined as set ofbodies or molecules divided from its surroundings by clearly defined boundaries and exchanging with its surroundings only mechanical work, heat and mass. Obviously, any adsorbate is an inhomogeneous phase as - by definition - its molecules are subject to the surface forces of the sorbent atoms. Hence the conditions for local (x = (xi, X2, X3)) thermodynamic equilibrium conditions, derived from the Second Law of thermodynamics, are [1.64]... 

Thermodynamic description of natural processes usually begins by dividing the world into a system and its exterior, which is the rest of the world. This cannot be done, of course, when one is considering the thermodynamic nature of the entire universe. The definition of a thermodynamic system often depends on the existence of boundaries, boundaries that separate the system of interest from the rest of the world. In understanding the thermodynamic behavior of physical systems, the nature of the interaction between the system and the exterior is important. Accordingly, thermodynamic systems are classified into three types, isolated, closed and open systems according to the way they interact with their exterior (Fig. 1.1). 

---