1. Thermodynamics
Science of energy deals with conversion of energy from
one form to another.
Energy
The amount of work that can be performed by a force to
produce change.
The law of conservation of energy.
Energy can neither be created (produced) nor destroyed by
itself. It can only be transformed from one form to another.
According to conservation law the total inflow of energy
into a system must equal the total outflow of energy from
the system, plus the change in the energy contained within
the system

2. Thermodynamic system.
In thermodynamics, a thermodynamic system, originally
called a working substance, is defined as that part of the
universe that is under consideration. Anything under
consideration is called a system. A boundary separates
the system from the rest of universe, which referred to
as environment, surroundings, or reservoir. A useful
classification of thermodynamics systems is based on
the nature of the boundary and the quantities flowing
through it, such as matter, energy, work, heat.

3. Surroundings

4. Isolated system
In isolate system, any modification of the environment has no
effect on the system, and any change of the system has no effect
on the environment neither matter nor energy in any form are
exchanged between the system. An example of an isolated
system would be an insulated rigid container, such as an
insulated gas cylinder.
Closed system.
A closed system is a system which may exchange energy in any
form with environment ( work, heat,….) but which cannot
exchange any matter.
Open system.
An open system can exchange both energy and matter with its
environment. The ocean would be an example of an open
system.

5. Zeroth law of thermodynamics, about thermal
equilibrium.
If two thermodynamics system are separately in
thermal equilibrium with a third, they are also in
thermal equilibrium with each other.
Properties of system.
1- Pressure (p).
2- Temperature (T) or (t).
3- Volume (V).
4- Mass (m).
Some are defined in term of other one
Density : mass per unit volume. ρ=m/V = (Kg/ m3)
Relative density ρs = ρ/ρH2o dimension less quantity.
Specific volume: volume per unit mass .ν= V/m = 1/ρ
= m3/kg.

6. Thermodynamics coordinate.
(P, V, and T) or (P, ρ, and T ) are thermodynamic
coordinate any change of these three coordinate gave
thermodynamic process.
1- An isobaric process occurs at constant pressure.
2- An isochoric process occurs at constant volume.
3- An isothermal process occurs at constant
temperature.

7. Equilibrium
A state of balance. A system which is in
equilibrium expressing no change from it’s
surrounding. There are many types of equilibrium
•Two systems are in thermal equilibrium when
their temperatures are the same (no temp.
difference).
• Two system are in mechanical equilibrium when
their pressures are the same (no pressure change
with time).
• Two system in chemical equilibrium when their
chemical potentials are the same (no chemical
reaction occurs).

8. Adiabatic process.
In thermodynamics, an adiabatic process is a thermodynamic
process in which no heat is transferred to or from the system.(The
quantity of thermal energy is constant).
Q1=Q2
dQ=0 Q= constant.
Temperature scales.
C0= (F0 - 32)*5/9
C0= K - 273.15
F0 = R0 – 459.67
Note
ΔK0=ΔC0
ΔR0 = ΔF0
C0 = degree Celsius.
F0 = degrees Fahrenheit.
R0 = Rankine .
K= Kelvin.

9. Equation of state
an equation of state is a relation between state
variables describing the state of matter under
a given set of physical conditions. It is a
constitutive equation which provides
mathematical relationship between two or
more state functions associated with the
matter, such as its (T, P and V).

10. Boyle’s law.
When gas is kept at constant temperature its
pressure is inversely proportion to the volume.
P1 V1 = P2 V2
Charles’s law.
When the pressure of the gas kept constant the
volume directly proportional to the temperature.
V1/T1 = V2/T2
The General Law
P1V1/T1 = P2V2/T2
Combined Boyle’s law and Charles’s law into the
first statement of ideal gas law.
PV=nRT n=number of moles
R= gas constant= 8.314 (j/mole.K)
T = temp. In (K) .

11. Real gas equation of state
Real gas has distance between molecules
less than distance between ideal gas molecule.
In real gas we have two equation.
1- Glasius equation.
P(V- nb)=nRT
2- Vander Waals equation.
(P + a/V3)(V- b)= RT
Where a and b are constant for any one gas
but differ for different gasses.

12. Heat
The form of energy that is transferred between
two systems by temperature difference. Heat
transferred during the process between two states
(state1 and state 2) is denote by Q12 or Q. Heat
has energy unit KJ or J.
q=Q/m (KJ/Kg). heat transfer per unit mass.

13. WORK .
Work W is performed whenever a force acts through a distance.
By definition, the quantity of work is given by the equation:
Where F is the component of force acting along the line of the
displacement dl. When integrated, this equation yields the work
of a finite process. By convention, work is regarded as positive
when the displacement is in the same direction as the applied
force and negative when they are in opposite directions.

14. or, since A is constant,
Integrating,

15. Figure (**) shows apath for compression of a gas from point 1
with initial volume V1t at pressure P1 to point 2 with volume
Vt2 at pressure P2. This path relates the pressure at any point
of the process to the volume. The work required is given by
Eq. (*) and is proportional to the area under the curve of Fig.
(**). The unit of work is the newton-meter or joule, symbol J.

16. THE FIRST LAW OF THERMODYNAMICS
The recognition of heat and internal energy as forms of energy
makes possible a generalization of the law of conservation of
mechanical energy include heat and internal energy in addition to
work and external potential and kinetic energy. Indeed, the
generalization can be extended to still other forms, such as surface
energy, electrical energy, and magnetic energy. This generalization
was at first a postulate. However, the overwhelming evidence
accumulated over time has elevated it to the stature of a law of
nature, known as the first law of thermodynamics. One formal
statement is:
Although energy assumes many forms, the total quantity of
energy is constant, and when energy disappears in one form it
appears simultaneously in other forms.

17. In application of the first law to a given process, the sphere of
influence of the process is divided into two parts, the system and
its surroundings. The region in which the process occurs is set
apart as the system; everything with which the system interacts
is the surroundings. The system may be of any size depending
on the application, and its boundaries may be real or imaginary,
rigid or flexible. Frequently a system consists of a single
substance; in other cases it may be complex. In any event, the
equations of thermodynamics are written with reference to some
well-defined system. This focuses attention on the particular
process of interest and on the equipment and material directly
involved in the process. However, the first law applies to the
system and surroundings, and not to the system alone. In its
most basic form, the first law requires:

18. Where the difference operator "Δ" signifies finite changes in the
quantities enclosed in parentheses. The system may change in its
internal energy, in its potential or kinetic energy, and in the
potential or kinetic energy of its finite parts. Since attention is
focused on the system, the nature of energy changes in the
surroundings is not of interest.
In the thermodynamic sense, heat and work refer to energy in
transit across the boundary which divides the system from its
surroundings. These forms of energy are not stored, and are never
contained in a body or system. Energy is stored in its potential,
kinetic, and internal forms; these reside with material objects and
exist because of the position, configuration, and motion of matter.

19. The choice of signs used with Q and W depends on which
direction of transport is regarded as positive. Heat Q and work W
always refer to the system, and the modern sign convention makes
the numerical values of both quantities positive for transfer into
the system from the surroundings. The corresponding quantities
taken with reference to the surroundings, Qsurr and Wsurr havthe
opposite sign, i.e., Qsurr = - Q and Wsurr = - W. With this
understanding:
Closed systems often undergo processes that cause no change in the system
other than in its internal energy. For such processes, reduces to:

20. HEATCAPACITY
We remarked earlier that heat is often viewed in relation to its effect on the
object to which or from which it is transferred. This is the origin of the idea
that a body has a capacity for heat. The smaller the temperature change in a
body caused by the transfer of a given quantity of heat, the greater its
capacity. Indeed, a heat capacity might be defined:
The difficulty with this is that it makes C, like Q, a process-dependent
quantity rather than a state function. However, it does suggest the possibility
that more than one useful heat capacity might be defined. In fact two heat
capacities are in common use for homogeneous fluids; although their names
belie the fact, both are state functions, defined unambiguously in relation to
other state functions.

21. Heat Capacity at Constant Volume
Heat Capacity at Constant Pressure
Isothermal Process
Isobaric Process

22. lsochoric (Constant- V) Process