Terminology of Thermodynamics
The important parts of the study of thermodynamics are a few terms and definitions, which must be understood clearly, and these are as follows:
1. System, boundary and surroundings: A thermodynamic system may be defined as any specified portion of matter in the universe which is under study. A system may consist of one or more substances.
The rest of the universe which might be in a position to exchange energy and matter with the system is called the surroundings.
Thus, the system is separated from the surroundings by a boundary which may be real or imaginary.
In experimental work, a specific amount of one or more substances constitutes the system. Thus 200 g of water contained in a beaker constitutes the thermodynamic system. The beaker and the air in contact are the surroundings.
Similarly, 1 mole of oxygen confined in a cylinder fitted with a piston, is a thermodynamic system. The cylinder and the piston and all other objects outside the cylinder, form the surroundings. Here the boundary between the system (oxygen) and the surroundings (cylinder and the piston) are clearly defined.
2. Homogeneous and Heterogeneous system: A system is said to be homogeneous when it is completely uniform throughout, for example, a pure solid or liquid or a solution or a mixture of gases. In other words, a homogeneous system consists of only one phase‘.
A system is said to be heterogeneous when it is not uniform throughout. In other words, a heterogeneous system is one which consists of two or more phases.
Thus a system consisting of two or more immiscible liquids or a solid in contact with a liquid in which it does not dissolve, is a heterogeneous system. A liquid in contact with its vapour is also a heterogeneous system because it consists of two phases.
3. Types of Thermodynamic Systems: There are three types of thermodynamic systems, depending on the nature of the boundary which are as follows:
(i) Isolated system: When the boundary is both sealed and insulated, no interaction is possible with the surroundings. Therefore,
An isolated system is one that can transfer neither matter nor energy to and from, its surroundings. Consider an example of a system consisting of a liquid in contact with its vapour in a closed vessel. Since the vessel is closed, no matter (liquid vapours) can go out or enter in the vessel. If the vessel is insulated also (as shown into figure), it can neither lose nor gain heat from the surroundings. Thus this system is an isolated system. A substance, say boiling water, contained in a thermos flask, is another example of an isolated system.
(ii) Closed system: Here the boundary is sealed but not insulated. Therefore, A closed system is one which cannot transfer matter but can transfer energy in the form of heat, work and radiation to and from its surroundings.
A specific quantity of hot water contained in a sealed tube is an example of a closed system. While no water vapour can escape from this system, it can transfer heat through the walls of the tube to the surroundings.
A phase may be defined as a homogeneous and physically distinct part of a system which is bound by a surface and is mechanically separable from other parts of the system.
A gas contained in a cylinder with a piston constitutes a closed system. As the piston is raised, the gas expands and transfers heat (energy) in the form of to the surroundings.
(iii) Open system: In such system the boundary is open and unisolated therefore,
An open system is one which can transfer both energy and matter to and from its surroundings.
Hot water contained in a beaker on laboratory table is an open system. The water vapour (matter) and also heat (energy) is transferred to the surroundings through the imaginary boundary. Zinc granules reacting with dil. HCI to produce gas in a beaker is another example of open system. gas escapes and the heat of the reaction is transferred to the surroundings.
4. Macroscopic System: The word macroscopic means ‘on a large scale.’ This term, therefore, is used to convey the sense of appreciable quantities i.e. quantities which can be weighed.
A system is said to be macroscopic when it consists of a large number of molecules, atoms or ions.
5. Macroscopic Properties: The properties associated with a macroscopic system are called macroscopic properties.
These properties are pressure, volume, temperature, composition, density, viscosity, surface tension, refractive index, colour etc.
6. State of System and State Variables: When macroscopic properties of a system have definite values, the system is said to be in a definite state. Whenever there is a change in any one of the macroscopic properties, the system is said to change into a different state. Thus the state of a system is fixed by its macroscopic properties.
Since the state of a system changes with the change in any of the macroscopic properties, these are called state variables. It also follows that when a system changes from one state (called initial state) to another state (called final state), there is invariably a change in one or more of the macroscopic properties.
Pressure, temperature, volume, mass and composition are the most important variables. In actual practice it is not necessary to specify all the variables because some of them are interdependent. In the case of a single gas, composition is not one of the variables because it remains always 100%.
Further, if the gas is ideal and one mole of the gas is under examination, it obeys the gas equation, PV = RT, where R is the universal gas constant. Evidently, if only two of the three variables (P, V and T) are known, the third can be easily calculated. Let the two variables be temperature and pressure. These are called independent variables. The third variable, generally volume, is said to be a dependent variable as its value depends upon the values of P and T. Thus, the thermodynamic state of a system consisting of a single gaseous substance may be completely defined by specifying any two of the three variable e.g. temperature, pressure and volume.
In a closed system, consisting of one or more components mass is not a state variable.
7. Extensive and Intensive Properties: All macroscopic or bulk properties of the system (volume, pressure, mass etc.), irrespective of whether they are state variables or not can be divided into two classes:
(i) Extensive properties
(ii) Intensive properties.
An extensive property of a system is that which depends upon the amount of the substance or substances present in the system. The examples are mass, volume, energy, heat capacity, enthalpy, entropy, free change etc.
An intensive property of a system is that which is independent of the amount of the substance present in the syste1′n.The examples are temperature, pressure, density, viscosity, refractive index, surface tension and specific heat.
The extensive properties are additive, while intensive properties are not. Sometimes an extensive property e becomes an intensive property by specifying unit amount of the substance concerned. Mass and volume are extensive properties but mass per unit volume i.e. density becomes an intensive property of the substance.
8. Thermodynamic Equilibrium: A system in which the macroscopic properties do not undergo any change with time is said to be in thermodynamic equilibrium.
Suppose a system is heterogeneous, if it is in equilibrium. The macroscopic properties in the various phases remain unchanged with time.
Actually, the term thermodynamic equilibrium implies the existence of three kinds of equilibria system. These are:
(i) Thermal equilibrium
(ii) Mechanical equilibrium
(iii) Chemical equilibrium
A system is said to be in thermal equilibrium. if there is no flow of heat from one position of the system tb another. This is possible if the temperature remains the same throughout in all parts of the system.
A system is said to be-tin mechanical equilibrium if no mechanical work is done by one part of the system on another part of the system. This is possible if the pressure remains the same throughout in all parts of the system.
A system is said to be in chemical equilibrium if the composition of the various phases in the system remains the same throughout.
9. Process: Whenever the state of a system changes, it is said to have undergone a process. Thus a process may be defined as the operation by which a system changes from one state to another.
In a process at least one of the properties of the system changes and gives a path along which the variables, of the system change. A change in state of the system is always accompanied by a change in energy. Therefore, a process may also be defined as a path of change of a system from one equilibrium state to another which is usually accompanied by a change in energy or mass.
Different types of processes connecting an initial state, in such changes, are discussed below: ‘
(i) Isothermal process (T remains constant): It is the process in which the temperature of the system remains constant during each step. In such a process the systems are in thermal contact with a constant temperature and both exchange heat with surroundings i.e. both maintain this temperature (T = 0).
(ii) Adiabatic process (Thermally insulated from the surroundings): A process in which no heat is exchanged between the system and surroundings is called adiabatic process (Q= O). System in which such processes occur are thermally insulated from the surroundings.
In such processes, the temperature of the system may change according to the conditions. For example, if heat is evolved in the system, the temperature of the system increases and if heat is absorbed, the temperature decreases.
(iii) process (V remains constant): A process in which the volume of the system remains constant during each step of the change in he state of the system is called isochoric process .
The reaction occurring in sealed containers of constant volume correspond to such processes.
(iv) Isobaric process (P remains constant): It is the process in which the pressure of the system remains constant during each step of the system ).
When a reaction occurs in an open beaker which will be at one atmosphere pressure, the process is called isobaric process.
Graphical representation of the process:
(v) Cyclic process: The process which brings aback a system to its original state after a series of changes is called a cyclic process.
Here = 0 (Enthaply change)
(Internal energy change)
(I) Isothermal expansion
(II) Adiabatic expansion
(III) Isothermic compression
(IV) Adiabatic compression
10. Reversible and Irreversible Processes: A thermodynamic reversible process is one that takes place infinitesimally slowly and its direction at any point can be reversed by an infinitesimally change in the state of the system.
In fact, a reversible process is considered to proceed from the initial state to the final state through an infinite series of infinitesimally small changes. At the initial, final and all intermediate stages, the system is in equilibrium state. This is so because an infinitesimal change in the state of the system at each intermediate step is negligible.
When a process goes from the initial state to the final state in a single step and cannot be carried in the reverse order, it is to be an irreversible process.
Here the system is in equilibrium state in the beginning and at the end, but not at points; in between; most of the processes are irreversible in nature. Flow of heat from high temperature to low temperature, water flowing from -hill, expansion of gas from higher to lower pressure, use of cells or batteries directly are few example of irreversible process. They are also called spontaneous process.
Consider a certain quantity of a gas contained in a cylinder having a weightless and frictionless piston. The expansion of the gas can be carried out by two methods illustrated in the figure below.
(A) Reversible expansion occurs by Decreasing the pressure on the Piston by infinitesimal amount.
(B) Irreversible expansion occurs by sudden decrease of pressure P to P’, when the gas expands rapidly in a single operation.
Let the pressure applied to the piston be P and this is equal to the internal pressure of the gas. Since the external and internal pressures are exactly counter-balanced, the piston remains stationary and there is no change in volume of the gas. Now, suppose the pressure on the piston is decreased by an infinitesimal amount dP. Thus the external pressure on the piston being P-dP, the piston moves up and the gas will expand by an infinitesimally small amount. The gas will therefore, be expanded infinitely slowly i.e. by a thermodynamically reversible process. At all stages in the expansion of the gas, dP being negligibly small, the gas is maintained in a state of equilibrium throughout. If at any point of the process, the pressure is increased by dP, the gas would contract reversibly;
On the other hand, the expansion is irreversible (Fig). if the pressure on the piston is decreased suddenly. It moves upward rapidly in a single operation. The gas is in equilibrium state in the initial and final stages only. The expansion of the gas, case, takes place in an irreversible manner.
Differences between Reversible and Irreversible Processes:
(i) Since reversible process takes place by infinitesimal small steps, the process would take infinite time to occur. Such a process is-idealized and is true in principle only. On the other hand, an irreversible process takes place in finite time. Thus all processes which actually occur are irreversible.
(ii) A reversible process is in equilibrium state at all stages of the operation, while an irreversible process is ‘in equilibrium state only at the initial and final stages of the operation.
(iii) A reversible process is unreal as it assumes the presence of frictionless and weightless piston. An irreversible process is real and can be performed actually.
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