Definition of Thermodynamics:
Thermodynamics is the study of effects of work, heat, and energy on the system. It deals only with large-scale response of the system that can be observed and estimated in the experiment, of heat and work. Small-scale gas interactions are explained by kinetic theory of gases.
Idea of a System and its Surrounding:
System is the restricted region of space or the finite portion of matter one has chosen to study. Or the part of universe, with well-defined boundaries, one has chosen to study.
Surrounding is rest of universe outside the region of interest (that is the rest of space outside system).
Boundary or Wall is surface which divides system from surroundings. This wall or boundary may or may not permit interaction between system and surroundings.
These are macroscopic coordinates or properties utilized to explain or characterize the system. As they are macroscopic properties or coordinates, they can be seen and estimated. Some examples are Temperature (T), thermal conductivity ( k ), Volume (V), specific heat capacity at constant pressure (CP ), Pressure (P), density (Ρ), mass (m), specific heat capacity at constant volume (CV ), thermal diffusivity (α), and chemical potential (μ).
This is a system which could be explained in terms of thermodynamic coordinates or properties. Thermodynamic Systems can be categorized into followings depending on kind of boundary:
Open System: This is a system that its boundary permits transfer of mass and energy into or out of system. In other words, boundary permits exchange of mass and energy between system and surrounding.
Closed System: This is a system that its boundary lets exchange of energy alone (inform of heat) between system and its surrounding (i.e. boundary lets exchange of energy alone). This kind of boundary which lets exchange of heat is known as diathermal boundary.
Isolated System: This is a system that its boundary lets neither mass nor energy between it and surrounding. In other words, boundary doesn't permit exchange of mass nor energy.
There are numerous specific kinds of thermodynamic processes which take place often enough (and in practical situations) that they are generally treated in study of thermodynamics. Each has unique feature which identifies it, and that is helpful in examining energy and work change related to process.
Adiabatic process: This is the thermodynamic procedure in which there is no heat transfer into or out of system. For this procedure, change in quantity of heat is zero (i.e. ΔQ = 0 during this process)
Isochoric process: This is the thermodynamic procedure which takes place at constant volume (i.e. ΔV = 0 during this procedure). This implies that during process no work is done on or by system.
Isobaric process: This is the thermodynamic process which takes place at constant pressure (i.e. Δp=0 during this procedure).
Isothermal process: This is the thermodynamic procedure which occurs at constant temperature (i.e. ΔT = 0 during this procedure) It is possible to have multiple procedures within the single process. A good example would be the case where volume and pressure change during the process, resulting in no change in temperature and no heat transfer. This type of a procedure would be both adiabatic and isothermal.
Cyclic Processes: These are series of procedures in which after certain interchanges of heat and work, system is restored to initial state. For the cyclic process ΔU =0 , and if this is put in first law Q =W This means that net work done during this procedure should be exactly equal to net amount of energy transferred as heat; store of internal energy of system remains unchanged.
Reversible Process: The reversible process can be stated as one which direction can be reversed by the infinitesimal change in some properties of system.
Quasi-static Process: This is a process which is performed in such a way that at every instant, system departs only infinitesimal from equilibrium state (i.e. almost static). Therefore quasi-static process closely approximates the succession of equilibrium states.
Non-quasi-static Process: This is a process which is performed in such a way that at every moment, there is finite departure of the system from the equilibrium state.
Irreversible Process: The irreversible process can be stated as one which direction can't be reversed by the infinitesimal change in some properties of system.
Usually, a system is said to be in equilibrium when its properties don't change noticeably with time over interval of interest (i.e. observation time). The system is said to be in thermodynamic equilibrium with its surrounding or with another system if and only if system is in thermal equilibrium, in chemical equilibrium and in mechanical equilibrium with surrounding or with another system. If any one of the above states is not fulfilled, system is not in thermodynamic equilibrium.
The system reaches mechanical equilibrium with its surrounding or with another system when there is no unbalance or net force in interior of the system and also none between system and its surroundings or another system. Assume two systems are separated by the movable boundary which doesn't permit exchange of mass or heat. If P1 is greater than P2, partition will continue to move toward system 2 until P1 is equal to P2. When this takes place, two systems are said to be in mechanical equilibrium.
The system achieves chemical equilibrium when there are no chemical reactions going on inside system or there is no transfer of matter from one part of system to other because of diffusion. Two systems are said to be in chemical equilibrium with each other when chemical potentials are same.
This takes place when two systems in thermal contact or the system which is in thermal contact with surrounding achieves same temperature. For instance if system 1 with temperature T1 and system 2 with temperature T2 are in thermal contact, there will be exchange of heat between two systems if there is temperature gradient (i.e. when T1 ≠ T2 ). This procedure of heat exchange will continue until thermal equilibrium is achieved (i.e. T1 = T2).
State of a System:
This is a specific situation in which macroscopic properties (thermodynamic properties) of a system have certain values (e.g. P=10 Pa, V=100 cm3, and T=300 K would be a state of a gas). It is important to note that the state of a pure substance or a system can be defined or specified by any two of its properties.
Change of state occurs when there is change in one, two or all the properties of the system. Using figure 1.1 above as example, suppose P1 is greater than P2 the partition will continue to move towards system 2 until P1 is equal to P2 . When this happens, the system 1 and 2 have a new set of coordinates in which Temperature remain constant for the two systems but pressure and volume changed. Then we say that the state of system 1 and system 2 has changed.
Equation of State:
This is the known relationship between thermodynamic variables or properties. It is the equation that gives the mathematical relationship between two or more state functions related with matter like its pressure, temperature, volume, or internal energy. Boyle's law, Charles' law, Dalton's law of partial pressures is examples of equation of state. Few other examples of equation of state are:
The Ideal Gas:
Equation of state for ideal gas is
PV = nRT
Where P is pressure, V is volume, R is molar gas constant (R=8.314 JK-1mol-1), T is temperature in Kelvin, and n is number of mole of gas.
Van der Waals Equation of State:
Equation of state for real gas also called as Van der Waals Equation is
(P + a/Vm)(Vm - b) = RT
Where quantities a and b are constants for the particular gas but vary for different gases.
Extensive and Intensive Properties:
Thermodynamic properties of the system can be categorized into two that is:
Extensive properties: These are properties of system which depend on mass of system (e.g. n, V and total energy U)
Intensive properties: These are properties of system which are independent of the mass of system (e.g. T, P and Ρ).
Specific Value of the extensive property (for example Volume, V) is stated as ratio of volume of property to mass of system, or as volume per unit mass.
Specific volume Vs is
Vs = V/m
Specific volume is clearly reciprocal of density ρ, stated as mass per unit volume:
Ρ = m/V = 1/Vs
Molar Value of the extensive property (for example Volume, V) is stated as ratio of volume of property to number of moles of system, or as volume per unit mole.
Molar volume Vm is:
Vm = V/n
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