1. Chapter Two: Basic Concepts of Thermodynamics_ Part One 1

2. System separates the system from its surroundings.
System is defined as a quantity of matter or a region in space chosen for study.
Boundary: The real or imaginary surface that
separates the system from its surroundings.
Surroundings: The mass or region outside the
system. 2 2

3. Closed System The volume of a closed system does not have to be fixed.
Closed system (also known as a control mass) consists of a fixed amount of mass, and no mass can cross its boundary (no mass can enter or leave the closed system).
Energy may cross the boundary of closed system
The volume of a closed system does not have to be fixed. 3

4. In isolated system, even energy is not allowed to cross the boundary4

5. Open System
In Open system (also called control volume) Both mass and energy can cross the boundary of a control volume.  
The mass of an open system does not have to be fixed. 5

6. System Properties
Any characteristic of a system is called a property. Some familiar properties are
Pressure (P)
Temperature (T)
Volume (V)
Mass (m)
Density (ρ), specific gravity or relative density
Specific volume (ν) 6

7. Types of Properties
Intensive properties are those that are independent of the mass of a system, such as temperature (T), pressure (P), and density (ρ).
Extensive properties are those whose values depend on the mass of the system such as : mass (M), volume (V), and total energy (E).
Specific properties. Some examples of specific properties are specific volume (v=V/m) and specific total energy (e=E/m).
The state of a simple compressible system is completely specified by two independent, intensive properties. 7

8. How To Differentiate Between Intensive and Extensive Properties?
• Divide the system into two equal parts with a partition
• If each part has the same property value as the
original system, the property is intensive
• If each part has only half the property value of the
original system, the property is extensive 8

9. State and Equilibrium
State: is the condition of a system at which a set of properties can completely describe it.
Equilibrium implies a state of balance.
Thermodynamics deals with equilibrium states. 9

10. Types of Equilibrium
Thermal Equilibrium: The temperature is the same throughout the entire system.
Mechanical Equilibrium: The pressure is the same throughout the entire system.
Phase Equilibrium: The mass of each phase reaches an equilibrium level and remains fixed.
Chemical Equilibrium: Chemical composition does not
change with time, i.e. no chemical reactions take place
Thermodynamic Equilibrium:
A system is said to be in thermodynamic equilibrium if it maintains thermal, mechanical, phase, and chemical equilibrium. 10 10

11. Processes, Paths, and Cycles
Process: Any change that a system undergoes from one equilibrium state to
another.
Path: The series of states through which a system passes during a process.
In order to describe a process completely, one should specify the initial and
final states AND the path the process follows.
Cycle: The system returns to its initial state at the end of the process. 11

12. Types of Processes
In most of the processes we study, one thermodynamic property is held constant.
The prefix iso- is often used to designate a process for which a particular property remains constant.
For examples:
Isothermal process: is a process during which the temperature T remains constant;
Isobaric process: is a process during which the pressure P remains constant.
Isochoric (or isometric) process: is a process during which the specific volume v remains constant.
Isentropic process: is a process during which the entropy remains constant. 12 12

13. Steady and Unsteady States
Steady state: implies no change with time.
Unsteady state: or transient, the opposite of steady state (there is change with time).
The steady-flow process:
a process during which a fluid flows through a control volume steadily (mass, volume and total energy content of the control volume remain constant). However, the fluid properties can change from point to point within the control volume, but at any fixed point they remain the same during the entire process 13 13

14. Forms of Energy
The total energy of a system is divided into two forms:
1. Macroscopic forms of energy are those which a system possesses as a whole with respect to some outside reference frame, such as kinetic and potential energies .
2. Microscopic forms of energy are those which are related to the molecular structure of a system and the degree of the molecular activity. The sum of all microscopic forms of energy is called the internal energy of a system and is denoted by U. 14

15. Forms of Energy
Kinetic energy: The energy that a system possesses as a result of its motion relative to some reference:
KE = ½ m v2 (J)
Potential energy: The energy that a system possesses as a result of its elevation in a gravitational field:
PE = mgh (J) 15 15

16. Forms of Energy
The total energy of a system consists of the kinetic, potential, internal energies, etc (i.e., the sum of all form of energy):
E = U + KE + PE
= U + ½ m v2 + mgh
Furthermore, most closed systems remain stationary (no change in their kinetic and potential energies). Therefore, for stationary systems
E = U 16 16

17. Types of associated energies with internal U
Portion of the internal energy of a system associated with the kinetic energies of the molecules is called the sensible energy.
The internal energy associated with the phase of a system is called the latent energy.
The internal energy associated with the atomic bonds in a molecule is called chemical energy.
The energy associated with the strong bonds within the nucleus of the atom itself is called nuclear energy. 17 17

18. Energy and Environment
The largest source of air pollution is the motor vehicles
The pollutants released by the vehicles are usually grouped as hydrocarbons (HC), nitrogen oxides (NOx), and carbon monoxide (CO).
There is also sulfur oxides (SOx) emitted to air from some Chemical Industries.
Some of the air pollutants cause Acidic Rain and Green House Effect. 18

19. Temperature and the Zeroth Law of Thermodynamics:
when an object is in contact with another object that is at a different temperature, heat is transferred from the object at higher temperature to the one at lower temperature until both objects attain the same temperature (Thermal Equilibrium).
The zeroth law can be restated as two objects are in thermal equilibrium if both have the same temperature reading even if they are not in contact. 19

20. Absolute Temperature P Experimental data Gas (A) Gas (B) extrapolation
extrapolation
Experimental data
Gas (A)
Gas (B)
Gas (C)
Gas (D) P
T ( C ) 20

21. Temperature We use four different measures of temperature:
Celsius (°C), Fahrenheit (°F), Kelvin (K), Rankine (R).
Kelvin and Rankine are absolute temperature scales. Both are zero at absolute zero K=0 R
Celsius and Fahrenheit are relative temperature scales, and
0°C= K °C=32°F
0°F=460 R °C=212°F 21

22. Temperature Conversion
Celcius
kelvin
1 DoC = 1 D K
1 DoC = (1.8) DoF
Fahrenheit
Rankine
1 DoF = 1 DR
1 DK = (1.8) DR
In ordinary usage the D symbol for the unit temperature difference is omitted. Consequently, you have to infer from the text whether oF, oC, oR, and K mean temperature or unit temperature difference. 22

23. Temperature Conversion
460 23

24. Example 1
Solution 24

25. Example 2
Convert 18 oC to oF
Convert 550 R to oC 25