An Introduction To Basic Laws Of Thermodynamic
Introduction
The four laws of Thermodynamics summarize the most important facts of thermodynamics. They define fundamental physical quantities such as temperature, energy and entropy, in order to describe thermodynamic systems. They also describe the transfer of energy as heat and work in thermodynamic processes.
1st Law of Thermodynamics
The 1st Law of Thermodynamics tells us that energy is neither created nor destroyed, thus the energy of the universe is a constant. However, energy can be transferred from one part of the universe to another. To work out .thermodynamic problems we will need to isolate a certain portion of the universe, the system, from the remainder of the universe, the surroundings.
The energy transfer between different systems can be expressed as:
where
E1 = initial energy
E2 = final energy
The internal energy encompasses:
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| FIG 1 |
The energy transfer between different systems can be expressed as:
E1 = E2
where
E1 = initial energy
E2 = final energy
The internal energy encompasses:
- The kinetic energy associated with the motions of the atoms
- The potential energy stored in the chemical bonds of the molecules
- The gravitational energy of the system
The first law is the starting point for the science of thermodynamics and for engineering analysis.
Based on the types of exchange that can take place we will define three types of systems:
- isolated systems: no exchange of matter or energy-
- closed systems: no exchange of matter but some exchange of energy-
- open systems: exchange of both matter and energy-
The first law makes use of the key concepts of internal energy, heat, and system work. It is used extensively in the discussion of heat engines.
Internal Energy
Internal energy is defined as the energy associated with the random, disordered motion of molecules. It is separated in scale from the macroscopic ordered energy associated with moving objects; it refers to the invisible microscopic energy on the atomic and molecular scale. For example, a room temperature glass of water sitting on a table has no apparent energy, either potential or kinetic . But on the microscopic scale it is a seething mass of high speed molecules. If the water were tossed across the room, this microscopic energy would not necessarily be changed when we superimpose an ordered large scale motion on the water as a whole.
Heat
Heat may be defined as energy in transit from a high temperature object to a lower temperature object. An object does not possess "heat"; the appropriate term for the microscopic energy in an object is internal energy. The internal energy may be increased by transferring energy to the object from a higher temperature (hotter) object - this is called heating.
- The potential energy stored in the chemical bonds of the molecules
- The gravitational energy of the system
The first law is the starting point for the science of thermodynamics and for engineering analysis.
Based on the types of exchange that can take place we will define three types of systems:
- isolated systems: no exchange of matter or energy-
- closed systems: no exchange of matter but some exchange of energy-
- open systems: exchange of both matter and energy-
The first law makes use of the key concepts of internal energy, heat, and system work. It is used extensively in the discussion of heat engines.
Internal Energy
Internal energy is defined as the energy associated with the random, disordered motion of molecules. It is separated in scale from the macroscopic ordered energy associated with moving objects; it refers to the invisible microscopic energy on the atomic and molecular scale. For example, a room temperature glass of water sitting on a table has no apparent energy, either potential or kinetic . But on the microscopic scale it is a seething mass of high speed molecules. If the water were tossed across the room, this microscopic energy would not necessarily be changed when we superimpose an ordered large scale motion on the water as a whole.
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| FIG 2 |
Heat may be defined as energy in transit from a high temperature object to a lower temperature object. An object does not possess "heat"; the appropriate term for the microscopic energy in an object is internal energy. The internal energy may be increased by transferring energy to the object from a higher temperature (hotter) object - this is called heating.
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| FIG 3 |
W = p dV
where
W is work,
p is pressure
dV is change in volume.
For non-constant pressure, the work can be visualized as the area under the pressure-volume curve which represents the process taking place.
For non-constant pressure, the work can be visualized as the area under the pressure-volume curve which represents the process taking place.
Heat Engines
Refrigerators, Heat pumps, Carnot cycle, Otto cycle
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| FIG 4 |
dU = Q - W
where
dU = change in internal energy
Q = heat added to the system
W = work done by the system
1st law does not provide the information of direction of processes and does not determine the final equilibrium state.
Intuitively, we know that energy flows from high temperature to low temperature. Thus, the 2nd law is needed to determine the direction of processes.
is the "thermodynamic potential" useful in the chemical thermodynamics of reactions and non-cyclic
processes. Enthalpy is defined by
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| FIG 5 |
H = U + PV
where
H = enthalpy
U = internal energy
P = pressure
V = volume
Entropy
is used to define the unavailable energy in a system. Entropy defines the relative ability of one system to act to an other. As things moves toward a lower energy level, where one is less able to act upon the surroundings,
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| FIG 6 |
the entropy is said to increase. Entropy is connected to the Second Law of Thermodynamics
For the universe as a whole the entropy is increasing.
2nd Law of Thermodynamics
he second law is concerned with entropy (S). Entropy is produced by all processes and associated with the entropy production is the loss of ability to do work. The second law says that the entropy of the universe increases. |
| FIG 7 |
For energy to be available there must be a region with high energy level and a region with low energy level. Useful work must be derived from the energy that would flows from the high level to the low level
- 100% of the energy can not be transformed to work
- Entropy can be produced but never destroyed
Entropy definition
Entropy is defined as :
S = H / T
where
S = entropy (kJ/kg K)
H = enthalpy (kJ/kg)
T = absolute temperature (K)
dS = dH / T
The sum of (H / T) values for each step in the Carnot cycle equals 0. This only happens because for every positive H there is a countering negative H, overall.
For a given physical process, the combined entropy of the system and the environment remains a constant if the process can be reversed. If we denote the initial and final states of the system by "i" and "f":
Sf = Si (reversible process)
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| FIG 8 |
An example of a reversible process is ideally forcing a flow through a constricted pipe. Ideal means no boundary layer losses. As the flow moves through the constriction, the pressure, temperature and velocity change, but these variables return to their original values downstream of the constriction. The state of the gas returns to its original conditions and the change of entropy of the system is zero. Engineers call such a process an isentropic process. Isentropic means constant entropy.
The second law states that if the physical process is irreversible, the combined entropy of the system and the environment must increase. The final entropy must be greater than the initial entropy for an irreversible process:
Sf > Si (irreversible process)
An example of an irreversible process is the problem discussed in the second paragraph. A hot object is put in contact with a cold object. Eventually, they both achieve the same equilibrium temperature. If we then separate the objects they remain at the equilibrium temperature and do not naturally return to their original temperatures. The process of bringing them to the same temperature is irreversible.
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| FIG 9 |
The efficiency of a heat machine working between two energy levels is defined in terms of absolute temperature:
where
η = efficiency
Th = temperature high level (K)
Tc = temperature low level (K)
As a consequence, to attain maximum efficiency the Tc would have to be as cold as possible. For 100% efficiency the Tc would have to equal 0 K. This is practically impossible, so the efficiency is always less than 1 (less than 100%).
Change in entropy > 0 Change in entropy = 0 Change in entropy < 0
irreversible process reversible process impossible process
η = ( Th - Tc ) / Th = 1 - Tc / Th
where
η = efficiency
Th = temperature high level (K)
Tc = temperature low level (K)
As a consequence, to attain maximum efficiency the Tc would have to be as cold as possible. For 100% efficiency the Tc would have to equal 0 K. This is practically impossible, so the efficiency is always less than 1 (less than 100%).
Change in entropy > 0 Change in entropy = 0 Change in entropy < 0
irreversible process reversible process impossible process
Entropy is used to define the unavailable energy in a system. Entropy defines the relative ability of one system to act on an other. As things move toward a lower energy level, where one is less able to act upon the surroundings, the entropy is said to increase.
For the universe as a whole the entropy is increasing!
Third Law of Thermodynamics
The entropy of a substance is zero if the absolute temperature is zero-the entropy of any pure substance in thermodynamic equilibrium approaches zero as the temperature approaches zero (Kelvin), or conversely
-the temperature (Kelvin) of any pure substance in thermodynamic equilibrium approaches zero when the entropy approaches zero
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| FIG 10 |
The Third Law of Thermodynamics can mathematically be expressed as
Nernst's Heat Theorem
lim S→0 = 0
T→0
whereS = entropy (J/K)
T = absolute temperature (K)
Since the absolute temperature zero can not be reached, no negative temperatures are possible.
The fact that there is a lowest temperature that one can achieve suggests that a convenient temperature scale is one that sets this temperature to zero
The zeroth law of thermodynamics
The zeroth law of thermodynamics is one of the four laws of thermodynamics, which states that if two systems are in thermal equilibrium with a third system, then they are in thermal equilibrium with one another. , these three objects will approach the same temperature.
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| FIG 11 |
Applications of the Zeroth Law of Thermodynamics
The thermometer may be the most well-known example of the zeroth law in action. For example, say the thermostat in your bedroom reads 67 degrees Fahrenheit. This means that the thermostat is in thermal equilibrium with your bedroom. However, because of the zeroth law of the thermodynamics, you can assume that both the room and other objects in the room (say, a clock hanging in the wall) are also at 67 degrees Fahrenheit.
Similar to the above example, if you take a glass of ice water and a glass of hot water and place them on the kitchen countertop for a few hours, they will eventually reach thermal equilibrium with the room, with all 3 reaching the same temperature.
If you place a package of meat in your freezer and leave it overnight, you assume that the meat has reached the same temperature as the freezer and the other items in the freezer.A Proposed Fifth Law of Thermodynamics
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