ELEMENTS OF MECH. ENGG. THERMODYNAMICS

1. ELEMENTS OF MECH. ENGG. THERMODYNAMICS Er.S.P.Singh Johal. Department of Mechanical Engineering PTU G.Z.S. Campus BATHINDA Phone9888111086 email :spsjohal@yahoo.com

2. Basic Concepts of Thermodynamics • Thermodynamics is the study of transformations of energy • System and surroundings • the system is the part of the world in which we have a special interest. A system has definite boundaries • the surroundings is everything outside the boundaries • Classification of systems: • open system can exchange matter as well as energy with its surroundings. • closed system can exchange energy with its surroundings. No transfer of matter across the boundaries is possible. • isolated system can exchange neither energy nor matter with its surroundings.

3. Department of Mechanical Engineering Applications of Thermodynamics Chap. 1: Basic Thermo. Concepts

4. Thermodynamics Systems • Thermodynamics system is defined as a quantity of matter or region in space chosen for study • The mass or region outside the system is called the surroundings • System boundary is the real and imaginary surface that separates the system from the surrounding. Boundary can be fixed or movable • May be closed or open

5. Open System / Control Volume • A system that involves mass and energy transfer across its boundaries is called an open system, or control volume • The boundaries of a control volume is called control boundaries and is fixed in shape and position • Energy in the form of heat and work as well as mass can cross the control boundaries

6. Open System / Control Volume Mass and Energy Cross Control Volume Boundaries

7. Closed System/Control Mass • A system of fixed mass is called a closed system, or control mass • The closed system boundary does not have to be fixed • No mass can cross the closed system boundary • Energy in the form of heat and work can cross the closed system boundary • If even energy is not allowed to cross we have an isolated system

8. Department of Mechanical Engineering Closed System/Control Mass Chap. 1: Basic Thermo. Concepts Energy, not mass, crosses closed-system boundaries Closed system with moving boundary

9. Work, Heat, and Energy • The energy of a system is a measure of its capacity to do work. • The energy of a system is the sum of the kinetic and potential energies of all particles in the system. • The energy of a closed system can be changed by: • work done on or by the system. • heat transfer across its boundaries. • Work is transfer of energy using organized motion. (expansion work, electrical work, etc.) • Heat is transfer of energy using thermal motion .(chaotic, random motion of molecules)

10. Heat Transfer • The boundary of a system is diathermic if heat can be transferred between system and surroundings. • The boundary of a system is adiabatic if heat cannot be transferred. • Adiabatic processes (no heat transfer between system and surroundings) take place in adiabatic systems. • A process that releases energy as heat is called exothermic. • A process that absorbs energy as heat is called endothermic.

11. Form of Energy • The sum of all forms of energy of a system is called Total Energy, which is considered to consist of internal, kinetic, and potential energies. E = U + mV2/2 + mgz • Internal energyrepresents the molecular energy of a system and may exist in sensible, latent, chemical, and nuclear forms. Represented by symbol, U. • Kinetic Energy is the energy that a system possesses as a results of its motion relative to some reference frame. KE = mV2/2 • Potential Energy is the energy that a system possesses as a results of its elevation in a gravitational field. PE = mgz

12. Internal Energy • The internal energy, U, is the total energy of a system. • We cannot give an absolute value of U but we can calculate U for a process. • U = Uf - Ui • Uf = final value of U • Ui = initial value of U • U is a state function .(the value of U depends only on the current state of the system) • U is an extensive property.

13. Properties of A System • Propertiesare any measurable characteristics of a system. eg. Pressure, temperature, volume, mass and density. • Extensive propertiesare the mass-dependent properties of a system. i.e. the properties that will vary proportionally with mass of the system. E.g. volume • Intensive propertiesare the properties that are not dependent on mass. Eg. Temperature, density. If any Extensive Property is divided by the mass we would also obtain an intensive property.

14. Intensive and Extensive Properties

15. State of a System • Definition - A set of properties that completely describe the conditions or characteristics of a system. • At a given state, all the properties of a system have fixed values. • State of a system will change when the properties of a system change.

16. Process, Path and Cycle • Process - Any change that a system undergoes from one equilibrium state to another is called a process. • Path- The series of state through which a system passes during a process is called a path. • Cycle- A process with identical end states is called a cycle.

17. State, Path, Process and Cycle Compressed Process P-V Diagram Each Point Along the Path is in Quasi-Equilibrium State If the Process returns to its initial State then we have a Cycle If the Outgoing and Returning Paths are Different ~ Net work is Produced (+ve or -ve)

18. Thermodynamic Equilibrium • Thermodynamics deals with Equilibrium States. • A system is said to be in thermodynamic equilibriumif it maintains thermal, mechanical, phase, and chemical equilibrium. • Thermal Equilibrium => Temperature is the same throughout the system. • Mechanical Equilibrium => Pressureis the same throughout the system. • Phase Equilibrium => No phase change process in the system. • Chemical Equilibrium => No chemical reactions

19. Quasi-Equilibrium Process • Definition- A process whereby the system remains infinitesimally close to an equilibrium states at all times. • During a quasi-static or quasi-equilibriumprocess, the system remains practically in equilibrium at all times. • A sufficiently slow process that allows the system to adjust itself internally so that properties in one part of the system do not change any faster than those at other parts.

20. Pressure • Pressure is defined as force per unit area. • Its unit is thepascal. • The absolute, gage, and vacuum pressures are related by.

21. Absolute, Vacuum and Gauge Pressure

22. Pressure Measurements • Small to moderate pressure differences are measured by a manometer, and a differential fluid column of height h corresponds to a pressure difference of . • The atmospheric pressure is measured by a barometer and is determined from.

23. Temperature and Zeroth Law of Thermodynamics • Temperature is a measure of ‘hotness’ or ‘coldness.’ • The zeroth law of thermodynamics states that two bodies are in thermal equilibrium with the third body then they are in equilibrium with each other. • Basis for validity of Temperature Measurement. • More fundamental than 1st and 2nd Laws of Thermodynamics.

24. Reversible Process • A process is regarded as thermodynamically reversible if it can be caused to go in either direction by an infinitesimal change in an external variable such as pressure or temperature. • Reversible changes occur when a system is in equilibrium with its surroundings. • For a reversible expansion: p = pex + dp • dp  0 • p = pex • w =  p dV

25. Indicator Diagram or PV-diagram • The expansion work, w, can be obtained from an indicator diagram. (a plot of p versus V) • The amount of work done by the gas is given by area under curve. • The maximum work available for a system operating between specified initial and final states is obtained when the change takes place reversibly . ( pex = p)

26. Expansion Work • dw = pexdv • dw is the expansion work (pressure-volume work) when a system undergoes a change. • pex is the external pressure. • dV is the change in volume. • dw >0 (work is done by the system) when system expands. (dV>0) • w =  pexdV • integration from Vi to Vf when volume changes from Vi to Vf.

27. Expansion Work, cont • Free Expansion - no opposing force • pex = 0 (expansion into a vacuum) • w = 0 • Expansion against Constant Pressure. • pex is constant. • w = pexV (V is the volume change) • Isothermal Reversible Expansion of Perfect Gas. • pex = p (reversible expansion) • p = nRT/V (ideal gas) • w =  nRT dV/V • w = nRT ln(Vf/Vi)

28. U for Process at Constant Volume • dq = du + dw • dw = dwexp + dwe • wexp is expansion work. (pressure-volume work) • we is other work . (electrical work etc.) • dwexp = 0 for a process taking place at constant volume. • dU = dq (if no electrical work) • dU = dqv (subscript v indicates process at constant volume) • U = qv for process at constant volume .

29. Heat Capacity at Constant Volume • CV = (U/T)V • CV is the heat capacity at constant volume. • (U/T)V is a partial derivative which shows how U varies with T when the volume is kept constant. • CV,m = CV/n • CV,m is the molar heat capacity at constant volume.

30. Change in U with T • dU = CV dT at constant volume • from definition of CV • U = CVT at constant volume • assuming that CV is independent of T • U = qv at constant volume • qv = CV T • qv is heat needed to change temperature by T

31. Enthalpy, H • H = U + pV (Definition of enthalpy) • H is a state function (U, p, and V are all state functions) • H is an extensive property • H = qp for a process taking place at constant pressure • assuming pressure-volume work is the only type of work involved in the process

32. Relation between H and U • H = U + pV • H = U + (pV) for a process (change) • (pV) is small for processes involving condensed phases (solids and liquids) only. • (pV) is generally significant for processes involving gases • H  U for processed involving condensed phases only. • (pV) = (nRT) for a gas (ideal gas) • H = U + (ngasRT) for processes involving gases.

33. The First Law • q = du + dw for a closed system • q = heat supplied to or removed from the system. (q <0 if heat removed from system) • w = work done on or by the system. • (w >0 if work done by the system) • q and w depend on the process by which the state is changed; they are not state functions. • U = 0 for an isolated system • the internal energy of an isolated system is constant. • dq = dw + du ( the First Law written for infinitesimal changes)