Short Notes: Zeroth and First Laws of Thermodynamics | Short Notes for Mechanical Engineering PDF Download

 Page 1


Zeroth and First Laws of Thermodynamics  
Zeroth Law of thermodynamics
• When two bodies have equality of temperature with a third body, then they 
have equality of temperature
• Definition of the zeroth law enables the use of a thermometer as a 
measurement device.
• A scale, however, needs to be defined. 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
• These quantities varied with pressure, however, so that different values would 
be obtained on top of a mountain versus down in the valley, and so this is not 
a good standard.
• The modern Celsius scale is defined to be nearly the same but has 0.01 °C as 
the so-called triple point of water, and-273.15 °C as absolute zero in K.
• The triple point of water is defined at the state where three phase of water 
(solid, liquid, and gas) are observed to co-exist.
• The transformation between the absolute Kelvin scale and the Celsius scale is 
given by
K =°C + 273.15
• The English equivalents are degrees Fahrenheit (°F) for relative temperature 
and degrees Rankine (°R) for absolute temperature. The conversions are
T(°R) = 1.8T(K),
T(°F) = 1,8T(°C) + 32, T(°F) = T(°R) - 459.67
Page 2


Zeroth and First Laws of Thermodynamics  
Zeroth Law of thermodynamics
• When two bodies have equality of temperature with a third body, then they 
have equality of temperature
• Definition of the zeroth law enables the use of a thermometer as a 
measurement device.
• A scale, however, needs to be defined. 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
• These quantities varied with pressure, however, so that different values would 
be obtained on top of a mountain versus down in the valley, and so this is not 
a good standard.
• The modern Celsius scale is defined to be nearly the same but has 0.01 °C as 
the so-called triple point of water, and-273.15 °C as absolute zero in K.
• The triple point of water is defined at the state where three phase of water 
(solid, liquid, and gas) are observed to co-exist.
• The transformation between the absolute Kelvin scale and the Celsius scale is 
given by
K =°C + 273.15
• The English equivalents are degrees Fahrenheit (°F) for relative temperature 
and degrees Rankine (°R) for absolute temperature. The conversions are
T(°R) = 1.8T(K),
T(°F) = 1,8T(°C) + 32, T(°F) = T(°R) - 459.67
Fixed Point Temp.(c)
Ice Point 0.01
Steam Point 100
Solidification of antimony 630.74
Solidification of Gold 1064.43
The First Law of Thermodynamics
• The first law of thermodynamics can be simply stated as follows: during an 
interaction between a system and its surroundings, the amount of energy 
gained by the system must be exactly equal to the amount of energy lost by 
the surroundings.
• A closed system can exchange energy with its surroundings through heat and 
work transfer. In other words, work and heat are the forms that energy can be 
transferred across the system boundary.
• Based on kinetic theory, heat is defined as the energy associated with the 
random motions of atoms and molecules.
• Heat is a directional (or vector) quantity; thus, it has magnitude, direction and 
point of action.
o Q (kJ) amount of heat transfer 
° Q° (kW) rate of heat transfer (power) 
o q (kJ/kg) - heat transfer per unit mass 
° q° (kW/kg) - power per unit mass
° Sign convention: Heat Transfer to a system is positive, and heat transfer 
from a system is negative. It means any heat transfer that increases the 
energy of a system is positive, and
heat transfer that decreases the energy of a system is negative
• Work is the energy interaction between a system and its surroundings. More 
specifically, work is the energy transfer associated with a force acting through 
a distance.
° W (kJ) amount of work transfer 
° W° (kW) power 
° w (kJ/kg) - work per unit mass 
° w° (kW/kg) - power per unit mass
° Sign convention: work done by a system is positive, and the work done 
on a system is negative.
Page 3


Zeroth and First Laws of Thermodynamics  
Zeroth Law of thermodynamics
• When two bodies have equality of temperature with a third body, then they 
have equality of temperature
• Definition of the zeroth law enables the use of a thermometer as a 
measurement device.
• A scale, however, needs to be defined. 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
• These quantities varied with pressure, however, so that different values would 
be obtained on top of a mountain versus down in the valley, and so this is not 
a good standard.
• The modern Celsius scale is defined to be nearly the same but has 0.01 °C as 
the so-called triple point of water, and-273.15 °C as absolute zero in K.
• The triple point of water is defined at the state where three phase of water 
(solid, liquid, and gas) are observed to co-exist.
• The transformation between the absolute Kelvin scale and the Celsius scale is 
given by
K =°C + 273.15
• The English equivalents are degrees Fahrenheit (°F) for relative temperature 
and degrees Rankine (°R) for absolute temperature. The conversions are
T(°R) = 1.8T(K),
T(°F) = 1,8T(°C) + 32, T(°F) = T(°R) - 459.67
Fixed Point Temp.(c)
Ice Point 0.01
Steam Point 100
Solidification of antimony 630.74
Solidification of Gold 1064.43
The First Law of Thermodynamics
• The first law of thermodynamics can be simply stated as follows: during an 
interaction between a system and its surroundings, the amount of energy 
gained by the system must be exactly equal to the amount of energy lost by 
the surroundings.
• A closed system can exchange energy with its surroundings through heat and 
work transfer. In other words, work and heat are the forms that energy can be 
transferred across the system boundary.
• Based on kinetic theory, heat is defined as the energy associated with the 
random motions of atoms and molecules.
• Heat is a directional (or vector) quantity; thus, it has magnitude, direction and 
point of action.
o Q (kJ) amount of heat transfer 
° Q° (kW) rate of heat transfer (power) 
o q (kJ/kg) - heat transfer per unit mass 
° q° (kW/kg) - power per unit mass
° Sign convention: Heat Transfer to a system is positive, and heat transfer 
from a system is negative. It means any heat transfer that increases the 
energy of a system is positive, and
heat transfer that decreases the energy of a system is negative
• Work is the energy interaction between a system and its surroundings. More 
specifically, work is the energy transfer associated with a force acting through 
a distance.
° W (kJ) amount of work transfer 
° W° (kW) power 
° w (kJ/kg) - work per unit mass 
° w° (kW/kg) - power per unit mass
° Sign convention: work done by a system is positive, and the work done 
on a system is negative.
Similarities between work and heat transfer
• Both are recognized at the boundaries of the system as they cross them 
(boundary phenomena).
• Systems posses energy, but not heat or work (transfer phenomena).
• Both are associated with a process, not a state. Heat or work has no meaning 
at a state.
• Both are path functions, their magnitudes depend on the path followed during 
a process as well as the end states
First Law of Thermodynamics
• First law, or the conservation of energy principle, states that energy can be 
neither created nor destroyed; it can only change forms.
• The first law cannot be proved mathematically, it is based on experimental 
observations, i.e., there are no process in the nature that violates the first law.
• The first law for a closed system or a fixed mass may be expressed as:
net energy transfer to (or from) the system as heat and work= net increase (or 
decrease) in the total energy of the system 
Q - W = AE (kJ)
where
Q = net heat transfer (=IQin - IQout)
W = net work done in all forms (=IWin - ZWout)
AE= net change in total energy (= E2 - E l)
• The change in total energy of a system during a process can be expressed as 
the sum of the changes in its internal, kinetic, and potential energies:
AE= AU + AKE + APE (kJ)
AU= m(u2-Ui)
AKE=1/2(mVz2- mV,2)
APE=mg(z2-z1 )
Note: for stationary systems APE=AKE=0, the first law reduces to
Q - W = AU
• The first law can be written on a unit-mass basis:
or in differential form:
q - w = Ae (kJ/kg)
5Q - 5W = dU (kJ)
Page 4


Zeroth and First Laws of Thermodynamics  
Zeroth Law of thermodynamics
• When two bodies have equality of temperature with a third body, then they 
have equality of temperature
• Definition of the zeroth law enables the use of a thermometer as a 
measurement device.
• A scale, however, needs to be defined. 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
• These quantities varied with pressure, however, so that different values would 
be obtained on top of a mountain versus down in the valley, and so this is not 
a good standard.
• The modern Celsius scale is defined to be nearly the same but has 0.01 °C as 
the so-called triple point of water, and-273.15 °C as absolute zero in K.
• The triple point of water is defined at the state where three phase of water 
(solid, liquid, and gas) are observed to co-exist.
• The transformation between the absolute Kelvin scale and the Celsius scale is 
given by
K =°C + 273.15
• The English equivalents are degrees Fahrenheit (°F) for relative temperature 
and degrees Rankine (°R) for absolute temperature. The conversions are
T(°R) = 1.8T(K),
T(°F) = 1,8T(°C) + 32, T(°F) = T(°R) - 459.67
Fixed Point Temp.(c)
Ice Point 0.01
Steam Point 100
Solidification of antimony 630.74
Solidification of Gold 1064.43
The First Law of Thermodynamics
• The first law of thermodynamics can be simply stated as follows: during an 
interaction between a system and its surroundings, the amount of energy 
gained by the system must be exactly equal to the amount of energy lost by 
the surroundings.
• A closed system can exchange energy with its surroundings through heat and 
work transfer. In other words, work and heat are the forms that energy can be 
transferred across the system boundary.
• Based on kinetic theory, heat is defined as the energy associated with the 
random motions of atoms and molecules.
• Heat is a directional (or vector) quantity; thus, it has magnitude, direction and 
point of action.
o Q (kJ) amount of heat transfer 
° Q° (kW) rate of heat transfer (power) 
o q (kJ/kg) - heat transfer per unit mass 
° q° (kW/kg) - power per unit mass
° Sign convention: Heat Transfer to a system is positive, and heat transfer 
from a system is negative. It means any heat transfer that increases the 
energy of a system is positive, and
heat transfer that decreases the energy of a system is negative
• Work is the energy interaction between a system and its surroundings. More 
specifically, work is the energy transfer associated with a force acting through 
a distance.
° W (kJ) amount of work transfer 
° W° (kW) power 
° w (kJ/kg) - work per unit mass 
° w° (kW/kg) - power per unit mass
° Sign convention: work done by a system is positive, and the work done 
on a system is negative.
Similarities between work and heat transfer
• Both are recognized at the boundaries of the system as they cross them 
(boundary phenomena).
• Systems posses energy, but not heat or work (transfer phenomena).
• Both are associated with a process, not a state. Heat or work has no meaning 
at a state.
• Both are path functions, their magnitudes depend on the path followed during 
a process as well as the end states
First Law of Thermodynamics
• First law, or the conservation of energy principle, states that energy can be 
neither created nor destroyed; it can only change forms.
• The first law cannot be proved mathematically, it is based on experimental 
observations, i.e., there are no process in the nature that violates the first law.
• The first law for a closed system or a fixed mass may be expressed as:
net energy transfer to (or from) the system as heat and work= net increase (or 
decrease) in the total energy of the system 
Q - W = AE (kJ)
where
Q = net heat transfer (=IQin - IQout)
W = net work done in all forms (=IWin - ZWout)
AE= net change in total energy (= E2 - E l)
• The change in total energy of a system during a process can be expressed as 
the sum of the changes in its internal, kinetic, and potential energies:
AE= AU + AKE + APE (kJ)
AU= m(u2-Ui)
AKE=1/2(mVz2- mV,2)
APE=mg(z2-z1 )
Note: for stationary systems APE=AKE=0, the first law reduces to
Q - W = AU
• The first law can be written on a unit-mass basis:
or in differential form:
q - w = Ae (kJ/kg)
5Q - 5W = dU (kJ)
For a cyclic process, the initial and final states are identical, thus AE=0. The 
first law becomes:
Q - W = 0 (kJ)
• Note: from the first law point of view, there is no difference between heat 
transfer and work, they are both energy interactions. But from the second law 
point of view, heat and work are very different.
Specific Heats
• The specific heat is defined as the energy required to raise the temperature of 
a unit mass of a substance by one degree. There are two kinds of specific 
heats:
° specific heat at constant volume, Cv (the energy required when the 
volume is maintained constant)
° specific heat at constant pressure, Cp (the energy required when the 
pressure is maintained constant)
• The specific heat at constant pressure Cp is always higher than Cv because at 
constant pressure the system is allowed to expand and energy for this 
expansion must also be supplied to the system.
• For a stationary closed system undergoing a constant-volume process (wb = 
0), Applying the first law in the differential form:
5q - 5w = du
at constant volume (no work) and by using the definition of Cv, one can write:
C d T = du
or
• Similarly, an expression for the specific heat at constant pressure Cp can be 
found. From the first law, for a constant pressure process (wb + Au = Ah). It 
yields:
• Specific heats (both Cv and Cp) are properties and therefore independent of 
the type of processes.
• Cv is related to the changes in internal energy u, and Cp to the changes in 
enthalpy,h.
• It would be more appropriate to define:
° Cv is the change in specific internal energy per unit change in 
temperature at constant volume.
° Cp is the change in specific enthalpy per unit change in temperature at 
constant pressure.
Specific heats for ideal gases
Page 5


Zeroth and First Laws of Thermodynamics  
Zeroth Law of thermodynamics
• When two bodies have equality of temperature with a third body, then they 
have equality of temperature
• Definition of the zeroth law enables the use of a thermometer as a 
measurement device.
• A scale, however, needs to be defined. 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
• These quantities varied with pressure, however, so that different values would 
be obtained on top of a mountain versus down in the valley, and so this is not 
a good standard.
• The modern Celsius scale is defined to be nearly the same but has 0.01 °C as 
the so-called triple point of water, and-273.15 °C as absolute zero in K.
• The triple point of water is defined at the state where three phase of water 
(solid, liquid, and gas) are observed to co-exist.
• The transformation between the absolute Kelvin scale and the Celsius scale is 
given by
K =°C + 273.15
• The English equivalents are degrees Fahrenheit (°F) for relative temperature 
and degrees Rankine (°R) for absolute temperature. The conversions are
T(°R) = 1.8T(K),
T(°F) = 1,8T(°C) + 32, T(°F) = T(°R) - 459.67
Fixed Point Temp.(c)
Ice Point 0.01
Steam Point 100
Solidification of antimony 630.74
Solidification of Gold 1064.43
The First Law of Thermodynamics
• The first law of thermodynamics can be simply stated as follows: during an 
interaction between a system and its surroundings, the amount of energy 
gained by the system must be exactly equal to the amount of energy lost by 
the surroundings.
• A closed system can exchange energy with its surroundings through heat and 
work transfer. In other words, work and heat are the forms that energy can be 
transferred across the system boundary.
• Based on kinetic theory, heat is defined as the energy associated with the 
random motions of atoms and molecules.
• Heat is a directional (or vector) quantity; thus, it has magnitude, direction and 
point of action.
o Q (kJ) amount of heat transfer 
° Q° (kW) rate of heat transfer (power) 
o q (kJ/kg) - heat transfer per unit mass 
° q° (kW/kg) - power per unit mass
° Sign convention: Heat Transfer to a system is positive, and heat transfer 
from a system is negative. It means any heat transfer that increases the 
energy of a system is positive, and
heat transfer that decreases the energy of a system is negative
• Work is the energy interaction between a system and its surroundings. More 
specifically, work is the energy transfer associated with a force acting through 
a distance.
° W (kJ) amount of work transfer 
° W° (kW) power 
° w (kJ/kg) - work per unit mass 
° w° (kW/kg) - power per unit mass
° Sign convention: work done by a system is positive, and the work done 
on a system is negative.
Similarities between work and heat transfer
• Both are recognized at the boundaries of the system as they cross them 
(boundary phenomena).
• Systems posses energy, but not heat or work (transfer phenomena).
• Both are associated with a process, not a state. Heat or work has no meaning 
at a state.
• Both are path functions, their magnitudes depend on the path followed during 
a process as well as the end states
First Law of Thermodynamics
• First law, or the conservation of energy principle, states that energy can be 
neither created nor destroyed; it can only change forms.
• The first law cannot be proved mathematically, it is based on experimental 
observations, i.e., there are no process in the nature that violates the first law.
• The first law for a closed system or a fixed mass may be expressed as:
net energy transfer to (or from) the system as heat and work= net increase (or 
decrease) in the total energy of the system 
Q - W = AE (kJ)
where
Q = net heat transfer (=IQin - IQout)
W = net work done in all forms (=IWin - ZWout)
AE= net change in total energy (= E2 - E l)
• The change in total energy of a system during a process can be expressed as 
the sum of the changes in its internal, kinetic, and potential energies:
AE= AU + AKE + APE (kJ)
AU= m(u2-Ui)
AKE=1/2(mVz2- mV,2)
APE=mg(z2-z1 )
Note: for stationary systems APE=AKE=0, the first law reduces to
Q - W = AU
• The first law can be written on a unit-mass basis:
or in differential form:
q - w = Ae (kJ/kg)
5Q - 5W = dU (kJ)
For a cyclic process, the initial and final states are identical, thus AE=0. The 
first law becomes:
Q - W = 0 (kJ)
• Note: from the first law point of view, there is no difference between heat 
transfer and work, they are both energy interactions. But from the second law 
point of view, heat and work are very different.
Specific Heats
• The specific heat is defined as the energy required to raise the temperature of 
a unit mass of a substance by one degree. There are two kinds of specific 
heats:
° specific heat at constant volume, Cv (the energy required when the 
volume is maintained constant)
° specific heat at constant pressure, Cp (the energy required when the 
pressure is maintained constant)
• The specific heat at constant pressure Cp is always higher than Cv because at 
constant pressure the system is allowed to expand and energy for this 
expansion must also be supplied to the system.
• For a stationary closed system undergoing a constant-volume process (wb = 
0), Applying the first law in the differential form:
5q - 5w = du
at constant volume (no work) and by using the definition of Cv, one can write:
C d T = du
or
• Similarly, an expression for the specific heat at constant pressure Cp can be 
found. From the first law, for a constant pressure process (wb + Au = Ah). It 
yields:
• Specific heats (both Cv and Cp) are properties and therefore independent of 
the type of processes.
• Cv is related to the changes in internal energy u, and Cp to the changes in 
enthalpy,h.
• It would be more appropriate to define:
° Cv is the change in specific internal energy per unit change in 
temperature at constant volume.
° Cp is the change in specific enthalpy per unit change in temperature at 
constant pressure.
Specific heats for ideal gases
It has been shown mathematically and experimentally that the internal energy 
is a function of temperature only
u=u(T)
• Using the definition of enthalpy (h = u + Pv) and the ideal gas equation of 
state (Pv = RT), we have: h = u + RT
• Since R is a constant and u is a function of T only: h = h(T)
• Therefore, at a given temperature, u, h, Cv and Cp of an ideal gas will have 
fixed values regardless of the specific volume or pressure.
• For an ideal gas, we have, du= CV (T) dT and dh= Cp(T) dT
• For an ideal gas, we can write:
R T = A (7’)-w(7’)
^ _ dh du 
~ dT dT 
R = c p-c,y
• The ratio of specific heats is called the specific heat ratio k = Cp/Cv
° varies with temperature, but this variation is very mild.
o for monatomic gases, its value is essentially constant at 1.67.
° Many diatomic gases, including air, have a specific heat ratio of about 
1.4 at room temperature.
Specific heats for solids and liquids
• A substance whose specific volume (or density) is constant is called 
incompressible substance.
• The specific volumes of solids and liquids (which can be assumed as 
incompressible substances) essentially remain constant during a process.
• The constant volume assumption means that the volume work (boundary 
work) is negligible compared with other forms of energy.
• As a result, it can be shown that the constant-volume and constant-pressure 
specific heats are identical for incompressible substances:
Cp = Cv = C
• For small temperature intervals, a C at averaged temperature can be used and 
treated as a constant, yielding:
AU= Cavg(T2-T-|)
• The enthalpy change of incompressible substance can be determined from 
the definition of enthalpy (h = u + Pv)
h2 - hi = (u2 - u l) + v(P2 - P I)
Ah = Au + vAP (kJ/kg)