Thermodynamic Concept: Internal Energy
Introduction:
With regards to Thermodynamics, the energy found within the system is referred to as Internal Energy. This energy is fundamental in the preparation of the system in any particular state. However, neither the kinetic energy (KE) nor potential energy (PE) as an entity is included. The Internal Energy tracks both the energy benefits and energy losses of the system that may results from any changes that may occur in its internal state. An increase in energy within a system can be achieved through the addition of matter, application of heat, and or through the incorporation of thermodynamic work. In instances where transfer of matter may be hindered by presence of impervious containing walls that system is referred to as a closed or confined system. The first law of thermodynamics, explains this fluctuation of internal energy as a total of the heat introduced into the system together with the thermodynamic work provided by the surrounding environment on the system. However, in cases where the containing walls don’t allow any form of matter or energy to pass, then the system is referred to as an isolated system that is characterized by its internal energy being rigid.
Definition of Important terms
Internal Energy of a given state: The quantification of energy of a given state within a system is challenging to achieve, and subject matter about the components is rarely fascinating. Thermodynamics mainly focuses on the adjustments associated with the internal energy, rather than the actual absolute value.
Changes of Internal energy: Changes of internal energy, related to standard state, are obtained from useful groups of thermodynamic activities and techniques from which a given state may be obtained.
Extensive Properties: The methods mentioned above can be characterized by specific broad variables of the system. These properties include mole numbers, entropy and electric dipole moment. For thermodynamics practice in general, it’s usually not possible to include all energies associated with the energy found within a system. Ordinarily, thermodynamic explanations center its focus on objects significant to the techniques being studied.
Internal energy is among the two principal objectives of the state variables. The amount of internal energy solely depends on the present state of the system but not on the procedure the system has experienced during its preparation. Internal energy is attributed to being an extensive quantity. It is the main thermodynamic potential, and all other thermodynamic potentials are usually derived from this type of energy. In the practice of thermodynamics it is not possible to acknowledge all forms of energy associated with internal energy. Energy within a system may also be described in atomic specifications through random kinetic energy and potential energy. Description through random kinetic energy results from movement of molecules within the system. Description of internal energy through potential energy is achieved through the presence of atomic forces arising due to the presence of atomic bonds. Internal energy usually comprises of both atomic kinetic and atomic potential energies Static rest mass-energy of the constituents of matter is fundamental for studies dealing with thermonuclear reactions.
UNIT of Internal Energy in various standards: Energy is usually expressed by the SI unit joule (J). However, for internal energy, an intensive energy density, referred to as specific internal energy, is often used. This is because this density can either be related to the mass of the system, units being J/kg, or be related to the amount of material having units in J/mol.
Description: The energy within a system U in a given state is usually calculated relative to a standard state of the system. This is achieved through the addition of the various macroscopic relocations of energy that guide the system from the standard state to the required state:
ΔU = Σ Ei
ΔU- denoting the difference in internal energy between the given state and standard state. And Σ Ei is the sum of energies that were conveyed into the system as it transitioned from the standard state to the given state.
The internal energy, may be classified into atomic potential energy, Umicro pot, and atomic kinetic energy, Umicro kin, components, from a microscopic point of view.
U = U microscopic potential + U microscopic kinetic
The microscopic kinetic energy results from movement of the machine’s particles relative to the center-of-mass frame. The microscopic potential energy comprises of energy resulting from chemical and nuclear molecular bonds, and energy as a result of deformation,
The energy found within a system does not encompass the energy due to the movement of molecules or position of molecules a system as a whole. Rather, the internal energy encompasses the input of such a field to the energy. This occurs often as a result of coupling between the subjective extent of movement of the object and the field. As a result, the field is added to the thermodynamic definition of the object taking the mode of an extra extraneous parameter.
In the practice of thermodynamics or engineering in general, it’s usually not possible to include all the various energies that make up the total inherent energy of a specimen system. In most cases, descriptions of sample systems include segments that are applicable to the sample under investigation. For most systems being investigated through thermodynamics, a handy null reference point is usually selected for internal energy.
The internal energy of a system is usually considered as an extensive attribute that is often influenced by the quantity of material in the system as well and the size of the system.
The atomic potential energy and kinetic energy are interchangeable at temperatures above the absolute zero. The total energy, however, remains the same in a confined system. Ordinarily, the kinetic energy usually diminishes at the point of zero temperature, and the total internal energy becomes purely comprised of the residual potential energy. This assumption is considered null as; quantum mechanics argues that molecules still possess some residual energy at zero temperature. This energy is commonly known as zero-point energy. Therefore any system at absolute zero is only at its lowest energy state and has attained the minimum energy it can possess.
Internal Energy of an IDEAL gas: The use of a perfect gas for learning reasons, and in the illustration of working systems is often incorporated in thermodynamics. By definition, an ideal gas usually consists of molecules that are considered as focal items that interact only by means of elastic interactions. These molecules tend to occupy a volume in such a manner that the path between interactions is greater than their diameter. Examples of such systems include noble gases such as helium and monatomic gases. For these gases, the kinetic energy is only made up of the translational energy of the separate atoms. For monoatomic molecules, there is neither rotation nor vibration of molecules. In addition, monoatomic molecules do not acquire excitement from electronically generated sources except in the presence of high temperatures.
Changes in internal energy changes of a perfect gas may be characterized only based on adjustments experienced in its kinetic energy of motion. The kinetic energy is affected by factors such as pressure, temperature and volume.
Summaries
Through this short and brief study of Internal energy. We came to know that what is the exact use of internal energy in our practical life. Such as where it is going to be used, what is its role in defining any particular thermodynamic system, what it possesses value at the particular situation of any system and what its dependency with the various parameter of the system. In addition, we were able to know what is its dependency with other thermodynamic parameters, and what is its unit, measurement, in relative sense or in an absolute sense when and where we should use, all these above-mentioned things about internal energy we came to know about after this short and brief research kind of review of this thermodynamic quantity.
References
Joule, J.P. (1850). “On the Mechanical Equivalent of Heat”. Philosophical Transactions
of the Royal Society. 140: 61–82. doi:10.1098/rstl.1850.0004.
Adkins, C.J. (1968/1975). Equilibrium Thermodynamics, second edition, McGraw-Hill,
London, ISBN 0-07-084057-1.
Bailyn, M. (1994). A Survey of Thermodynamics, American Institute of Physics Press,
New York, ISBN 0-88318-797-3
van Gool, W.; Bruggink, J.J.C. (Eds) (1985). Energy and time in the economic and
physical sciences. North-Holland. pp. 41–56. ISBN 978-0444877482.