Physics Program

1. THERMODYNAMICS
Department of Physics, Faculty of Science
Jazan University, KSA
Part-1
222 PHYS

2. Thermodynamics is the study of the effects of work, heat, and energy on a system
Thermodynamics is only concerned with macroscopic (large-scale) changes and
observations
Thermodynamics
• Thermodynamic system: a quantity of fixed mass under investigation,
Definitions
• System boundary: interface separating
system and surroundings
• Universe: combination of system and
surroundings.
• Surroundings: everything external to the system

3. Definitions
Systems can be classified as:
•Open: mass and energy can be transferred between system and surroundings
•Closed: energy can be transferred but not mass
•Isolated: Neither energy nor mass can be transferred between system and
surroundings
A Closed system (a controlled mass) consists of a fixed amount of mass, and
no mass can cross its boundary. That is, no mass enters or leave a
closed system.
such as, Piston-cylinder device
An Open system (or a control volume ） is a properly selected region in
space. Both mass and energy can cross the boundary of a control volume.
such as, A Water heater, a turbine and a compressor, etc

4. Definitions
Intensive properties: are those independent of the size of a system as
well as its Temperature (T), Pressure (P) and concentration and density
Extensive properties: are those whose values depend on the size or
extent of the system
Sum of the properties of the system’s components.
•Depends on the size of the system
•Volume (V), Area (A), number of moles (n), mass(m), internal energy (U)
and enthalpy (H)
Two important interactions between the system and its surroundings:
• heat can cross into the system and
• work can cross out of the system
All of thermodynamics can be expressed in terms of four quantities
Temperature (T) - Internal Energy (U) - Entropy (S) - Heat (Q)

5. Thermodynamics
provides a framework of relating the macroscopic properties of a system to one
another.
In general: Change where system is always in thermal equilibrium:
{ reversible process}
Change where system is not always in thermal equilibrium:
{irreversible process}
Examples of irreversible processes:
• Free expansion - melting of ice in warmer liquid - frictional heating
Types of the questions that Thermodynamics addresses:
• How does an engine work? What is its maximum efficiency?
• How does a refrigerator work? What is its maximum efficiency?
• How much energy do we need to add to a kettle of water to change it to steam?

6. Properties of a pure substance
Pure substancePure substance: a material with homogeneous and invariable composition.
• Pure substances can have multiple phases, It may exist in more than one phase,
but the chemical composition is the same in all phases.
•An ice-water mixture is still a pure substance.
• An air-steam mixture is not a pure substance.
(air, being composed of a mixture of N2, O2, and other gases,
is formally not a pure substance).

7. Phases of matter
Gas - very weak
intermolecular forces,
rapid random motion
Liquid - intermolecular
forces bind closest neighbours
Solid - strong
intermolecular forces
low
temp
high
pressure
high temp
low pressure

8. Property Definition Symbol S.I.Unit
Volume Volume of a substance V m3
Internal Energy
The translational, rotational and
vibrational kinetic energy of a
substance
U Joules (J)
Enthalpy U + PV H Joules (J)
Entropy
The entropy is a measure of the
lack of structure or the amount of
disorder in a system
S Joules/Kelvin (J / K)

9. Thermodynamics
•“set of tools” that describes the macroscopic properties of
equilibrium systems
•entirely empirical science
•based on four laws

10. The zeroth law of thermodynamics deals with thermal equilibrium and
provides a means of measuring temperature.
The first law of thermodynamics: deals with the conservation of
energy and introduces the concept of internal energy.
The second law of thermodynamics: dictates the limits on the
conversion of heat into work and provides the yard stick to measure
the performance of various processes. It also tells whether a
particular process is feasible or not and specifies the direction in
which a process will proceed. As a consequence it also introduces the
concept of entropy.
The third law of thermodynamics: defines the absolute zero of
entropy.

11. Zeroth law of thermodynamics
When two bodies have equality of temperature with a third
body,
then they have equality of temperature
If A and C are at thermal equilibrium, i.e. at the same temperature,
and B and C are at thermal equilibrium, then it follows that A and B
are at thermal equilibrium, i.e. at the same temperature.

12. Heat
Sign of Q : Q > 0 system gains thermal energy
Q < 0 system loses thermal energy
The term Heat (Q) is properly used to describe the thermal
energy transferred into or out of a system from a thermal
reservoir.
Q U
Heat is a transfer of energy.

13. Heat is a form of energy.
It is a scalar quantity.
It is measured in joules (J).
When we heat a substance it
becomes hotter
or changes state.

14. The old metric temperature scale, Celsius (◦C), was defined so that
0 ◦C is the freezing point of water, and
100 ◦C is the boiling point of water.
The transformation between temperature scales is as follows
Heat always flow from an object with excess energy to an object with an absence of energy.
Heat flows from hot to cold.
1 cal = 4.186 J
Convert calories to Joules
O
[ ] [ ] 273.15T K T C= + O 5
[ ] ( [ ] 32)
9
T C T F= −

15. The Absolute (Kelvin) Temp. Scale
The absolute (Kelvin) temperature
scale is based on fixing T of the triple
point for water (a specific T = 273.16 K
and P = 611.73 Pa where water can
coexist in the solid, liquid, and gas
phases in equilibrium).






≡
TPP
P
KT 16.273 - for an ideal gas constant-
volume thermoscope
absolute zero
T,K
PPTP
273.16
PTP – the pressure of the gas in a
constant-volume gas thermoscope
at T = 273.16 K
0

16. Heat CapacityHeat Capacity
• The heat capacity is the energy required to raise the temperature of an object by 1o
C.
• The heat capacity depends on the mass of the object and the ability of the object to
resist the flow of heat into or out of the object.
Q = mcxΔT
• The constant cx is called the specific heat of substance x, (SI units of J/kg·K)
1 dQ dq
c
m dT dT
 
≡ = ÷
 
Specific heat capacity
specific heat”, is the heat capacity per unit mass
m – mass [kg]
C – Heat capacity [J/o
C]
c – Specific heat [J/kgo
C]
0
lim
T
Q dQ
C
T dT∆ →
 
≡ = ÷
∆ 
Heat capacity

17. Because the differential dQ is inexact, we have to specify under what
conditions heat is added.
Or, more precisely, which parameters are held constant. This leads to two
important cases:
• the heat capacity at constant volume, CV
• the heat capacity at constant pressure, Cp
andV P
V P
dQ dQ
C C
dT dT
   
≡ ≡ ÷  ÷
   
( isothermic, C = ∞,
adiabatic, C = 0 )
3 5
Monatomic: ;
2 2
5 7
Diatomic: ; ;
2 2
V P
V P
R R
c c
R R
c c
 
= = 
 
 = =
  

18. A common use of the heat equation is to determine the final temperature of
a mixture of two different objects at different initial temperatures.
coldhot QQ −=
222111 TmcTmc ∆−=∆
This is necessary to
compensate for the different
directions of heat flow.
( ) ( )222111 TTmcTTmc ff −−=−
2222211111 TmcTmcTmcTmc ff +−=−
2221112211 TmcTmcTmcTmc ff +=+
( ) 2221112211 TmcTmcmcmcTf +=+
( )2211
222111
mcmc
TmcTmc
Tf
+
+
=
This would be the final temperature of a mixture of two
materials with different mass, specific heats and initial
temperatures.
This expression is not valid if there is a phase change at any time during the process.

19. Specific Heat values for selected substances.

20. Heat and Latent Heat
Latent heat is defined as the heat absorbed or released when a substance
changes its physical state completely at constant temperature
1) Latent heat fusion
2) Latent heat of vaporization
Latent heat fusion is the heat absorbed when a solid melt at constant
temperature
Latent heat of vaporization is the heat absorbed
when a liquid changes into vapor at constant
temperature

21. Latent Heat The amount of heat required for a phase transition is given by:
mLQ =
Q – Heat [J]
L – Latent Heat [J/kg]
M – mass [kg] Ckg
J
Cg
cal
cwater 
41841 ==
Ckg
J
Cg
cal
cice 
20925.0 ≈≈
Ckg
J
Cg
cal
csteam 
20105.0 ≈≈
kg
J
x
g
cal
L mf
5
/ 103.380 ==
kg
J
x
g
cal
L cv
6
/ 1026.2540 ==

22. Example1-1:
10 kg of ice at 0o
C is dropped into a steam chamber at 100o
C. How much steam is
converted to water at 100o
C?
hotcold QQ −=
2 2ice f H O H O water steam vm L c m T M L+ ∆ = −
2 2ice f H O H O water
steam
v
m L c m T
M
L
+ ∆
= −
kgMsteam 3.3−=
The negative sign on the mass is acceptable here since it
represents the mass of steam that was lost (converted to water).

23. Heat and Latent Heat
• Latent heat of transformation L is the energy required for 1
kg of substance to undergo a phase change. (J / kg)
Q = ±ML
• Specific heat c of a substance is the energy required to raise
the temperature of 1 kg by 1 K. (Units: J / K kg )
Q = M c ΔT
• Molar specific heat C of a gas at constant volume is the
energy required to raise the temperature of 1 mol by 1 K.
Q = n CV ΔT
If a phase transition involved then the heat transferred is
Q = ±ML+M c ΔT

24. Latent heat and specific heat
• The molar specific heat of gasses depends on the process path
• CV= molar specific heat at constant volume
• Cp= molar specific heat at constant pressure
Cp= CV+R (R is the universal gas constant)
VC
Cp
=γ