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Chemical Thermodynamics of Materials Macroscopic and Microscopic Aspects by Svein Stolen, Tor Grande, and Neil L. Allan | PDF Free Download.
A thermodynamic description of a process needs a well-defined system. A thermodynamic system contains everything of thermodynamic interest for a particular chemical process within a boundary.
The boundary is either a real or hypothetical enclosure or surface that confines the system and separates it from its surroundings.
In order to describe the thermodynamic behavior of a physical system, the interaction between the system and its surroundings must be understood.
Thermodynamic systems are thus classified into three main types according to the way they interact with the surroundings: isolated systems do not exchange energy or matter with their surroundings;
Closed systems exchange energy with the surroundings but not matter, and open systems exchange both energy and matter with their surroundings.
The system may be homogeneous or heterogeneous. An exact definition is difficult, but it is convenient to define a homogeneous system as one whose properties are the same in all parts, or at least their spatial variation is continuous.
A heterogeneous system consists of two or more distinct homogeneous regions or phases, which are separated from one another by surfaces of discontinuity.
The boundaries between phases are not strictly abrupt, but rather regions in which the properties change abruptly from the properties of one homogeneous phase to those of the other. For example, Portland cement consists of a mixture of the phases b-Ca2SiO4, Ca3SiO5, Ca3Al2O6, and Ca4Al2Fe2O10.
The different homogeneous phases are readily distinguished from each other macroscopically and the thermodynamics of the system can be treated based on the sum of the thermodynamics of every single homogeneous phase.
In colloids, on the other hand, the different phases are not easily distinguished macroscopically due to the small particle size that characterizes these systems.
So although a colloid also is a heterogeneous system, the effect of the surface thermodynamics must be taken into consideration in addition to the thermodynamics of each homogeneous phase.
In the following, when we speak about heterogeneous systems, it must be understood (if not stated otherwise) that the system is one in which each homogeneous phase is spatially sufficiently large to neglect surface energy contributions.
The contributions from surfaces become important in systems where the dimensions of the homogeneous regions are about 1 mm or less in size.
The thermodynamics of surfaces will be considered in Chapter 6. A homogeneous system – solid, liquid, or gas – is called a solution if the composition of the system can be varied.
The components of the solution are the substances of fixed composition that can be mixed in varying amounts to form the solution. The choice of the components is often arbitrary and depends on the purpose of the problem that is considered.
The solid solution LaCr1–yFeyO3 can be treated as a quasi-binary system with LaCrO3 and LaFeO3 as components.
Alternatively, the compound may be regarded as forming from La2O3, Fe2O3, and Cr2O3 or from the elements La, Fe, Cr, and O2 (g). In La2O3 or LaCrO3, for example, the elements are present in a definite ratio, and independent variation is not allowed.
La2O3 can thus be treated as a single component system. We will come back to this important topic in discussing the Gibbs phase rule in Chapter 4.
In thermodynamics, the state of a system is specified in terms of macroscopic state variables such as volume, V, temperature, T, pressure, p, and the number of moles of the chemical constituents I, ni.
The laws of thermodynamics are founded on the concepts of internal energy (U), and entropy (S), which are functions of the state variables.
Thermodynamic variables are categorized as intensive or extensive. Variables that are proportional to the size of the system (e.g. volume and internal energy) are called extensive variables, whereas variables that specify a property that is independent of the size of the system (e.g. temperature and pressure) are called intensive variables.
A state function is a property of a system that has a value that depends on the conditions (state) of the system and not on how the system has arrived at those conditions (the thermal history of the system).
For example, the temperature in a room at a given time does not depend on whether the room was heated up to that temperature or cooled down to it.
The difference in any state function is identical for every process that takes the system from the same given initial state to the same given final state: it is independent of the path or process connecting the two states.
Whereas the internal energy of a system is a state function, work and heat are not. Work and heat are not associated with one given state of the system but are defined only in a transformation of the system.
Hence the work performed and the heat absorbed by the system between the initial and final states depend on the choice of the transformation path linking these two states.
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