Thermal energy
The thermal energy or heat energy is the part of the internal energy of a thermodynamic system in equilibrium that is provided at its absolute temperature and is increased or decreased by transfer of energy, usually in the form of heat or work, in thermodynamic processes. At the microscopic level and within the framework of the kinetic theory, it is the total of the average kinetic energy present as a result of the random movements of atoms and molecules or thermal agitation, which disappear instantly.
Introduction
In 1807 Thomas Young coined the term energy and in 1852 Lord Kelvin proposed its use in thermodynamics. The concept of internal energy and its symbol U{displaystyle U} first appeared in the works of Rudolph Clausius and William Rankine, in the second half of the centuryXIX, and over time replaced the terms internal work, internal work and intrinsic energy used usually at that time. James Prescott Joule would introduce latent heat definitions and sensitive heat.
Thermal energy represents the total internal energy of an object: the sum of its potential and kinetic molecular energies. When two objects with different temperatures come into contact, energy is transferred from one to the other. For example, if hot coals are dropped into a container of water, heat energy will transfer from the coals to the water until the system reaches a stable condition called thermal equilibrium.
In thermodynamics, the thermal energy also known as the internal energy of a system is the sum of the kinetic energies of all its constituent particles, plus the sum of all the potential energies of interaction between them. Kinetic and potential energy are microscopic forms of energy, that is, they are related to the molecular structure of a system and the degree of molecular activity, and are independent of external reference frames; For this reason, it is important to clarify that the internal energy does not include the potential energy due to the interaction between the system and its environment, therefore, the internal energy of a substance does not include the energy that it may possess as a result of its macroscopic position or his move.
According to atomic theory, thermal energy represents the kinetic energy of fast-moving molecules. The rise in temperature corresponds to an increase in the average kinetic energy of the molecules. Because thermal energy represents the energy of the atoms and molecules that make up an object, it is often called internal energy. From an atomic point of view, internal energy can include not only the kinetic energy of molecules, but also potential energy (usually electrical in nature) due to the relative positions of atoms within molecules. At a macroscopic level, internal energy corresponds to non-conservative forces such as friction. At the atomic level, however, energy is partially kinetic and potential, and the corresponding forces are conservative.
The symbol is used U{displaystyle U} for internal energy. During a system state change, internal energy could change from an initial value U1{displaystyle U_{1}} to one final U2{displaystyle U_{2}}. Change in internal energy is noted as Δ Δ U=U2− − U1{displaystyle Delta U=U_{2}-U_{1}}}}.
When a certain amount of heat is added Q{displaystyle Q} to a system and this does not work during the process (so W=0{displaystyle W=0}), internal energy increases in an amount equal to Q{displaystyle Q}I mean, Δ Δ U=Q{displaystyle Delta U=Q}. When the system does a job W{displaystyle W} expanding against its environment and not adding heat during that process, energy comes out of the system and decreases internal energy: W{displaystyle W} It's positive. Q{displaystyle Q} is zero and this does not work during the process (so that W=0{displaystyle W=0}), internal energy increases in an amount equal to Q{displaystyle Q}I mean, Δ Δ U=− − W{displaystyle Delta U=-W}. If there is both heat transfer and work, the total internal energy change is:
U2− − U1=Δ Δ U=Q− − W{displaystyle U_{2}-U_{1}=Delta U=Q-W} (First Thermodynamic Law)
This can be rearranged as follows:
Q=Δ Δ U+W{displaystyle Q=Delta U+W}
This means when you add heat Q{displaystyle Q} to a system, a part of this added energy remains in the system, modifying its internal energy in a quantity Δ Δ U{displaystyle Delta U}the rest comes out of the system when it does a job W{displaystyle W} against their environment. Since W and Q can be positive, negative or zero, Δ Δ U{displaystyle Delta U} can be positive, negative or zero for different processes. The first law of thermodynamics is a generalization of the principle of energy conservation to include energy transfer as heat and as mechanical work.
Thermodynamic approach
In the thermodynamic analysis, it is often useful to consider two groups for the various forms of energy that make up the total energy of a system: macroscopic and microscopic. Macroscopic forms of energy are those that possess a system as a whole in relation to a certain external frame of reference, such as kinetic and potential energies. Microscopic forms of energy are those that relate to the molecular structure of a system and the degree of molecular activity, and are independent of external frames of reference. The sum of all microscopic forms of energy is called internal energy of a system and is denoted by U{displaystyle U}.
To better understand internal energy, systems are examined at the molecular level. Gas molecules move in space with a certain speed; therefore, they possess some kinetic energy. This is known as translational energy. The atoms of polyatomic molecules rotate about an axis, and the energy associated with this rotation is the kinetic energy of rotation. The atoms of this type of molecules could vibrate with respect to their common center of mass, then the energy of this “reciprocating” movement would be the vibrational kinetic energy. For gases, the kinetic energy is mainly due to translational and rotational motions, where the vibratory motion becomes significant at high temperatures. The electrons in an atom revolve around the nucleus and therefore have rotational kinetic energy. Electrons in outer orbits have larger kinetic energies. Since these particles also spin around their axes, the energy associated with this motion is spin (spin) energy. The other particles that are located in the nucleus of an atom also possess spin energy. The portion of the internal energy of a system related to the kinetic energy of the molecules is called sensible energy (or kinetic energy of molecules). The average velocity and the degree of activity of the molecules are proportional to the temperature of the gas, so at higher temperatures the molecules have higher kinetic energies and, as a consequence, the system has a higher internal energy. Internal energy is also related to various bonding forces between the molecules of a substance, between the atoms within a molecule, and between the particles within an atom and its nucleus. The forces that bind molecules together are, as would be expected, strongest in solids and weakest in gases. If enough energy is added to the molecules of a solid or liquid, they overcome the molecular forces and pull apart, so that the substance becomes a gas; this is a phase change process. Due to the added energy, a system in the gas phase is at a higher level of internal energy than that in the solid or liquid phase. The internal energy related to the phase of a system is called latent energy. The phase change process can occur without changing the chemical composition of a system. Most of the real problems fall into this category, so it is not necessary to pay attention to the bonding forces of the atoms in a molecule. An atom in its nucleus has positively charged neutrons and protons bound together by strong forces from heat flow, as well as negatively charged electrons orbiting around them. The internal energy related to the atomic bonds in a molecule is called chemical energy. During a chemical reaction, for example a combustion process, some chemical bonds are broken and others are formed, resulting in the internal energy undergoing a change. Nuclear forces are much greater than those that bind electrons to the nucleus. This enormous amount of energy related to the strong bonds within the nucleus of the atom is called nuclear energy.
Internal energy of an ideal gas
The kinetic energy of translation Ec{displaystyle E_{c}} of the molecules of an ideal gas is related to absolute temperature T.
Ec=32nRT{displaystyle E_{c}={frac {3}{2}}nRT}
where n is the number of moles of gas and R, the universal constant of gases. If this translation energy is considered to constitute all the internal energy of the gas then U=Ec{displaystyle U=E_{c}}:
U=32nRT{displaystyle U={frac {3}{2}nRT}
In this case the internal energy of an ideal gas depends only on its temperature and the number of moles, not on its pressure or volume. If molecules, in addition to the kinetic energy of translation, have other types of energy such as rotation energy, the internal energy will be greater than that expressed by the previous equation. However, according to the equipartition theorem, the average energy associated with any degree of freedom will be 12kT{displaystyle {frac {1}{2}}kT} by molecule or 12RT{displaystyle {frac {1}{2}}{2}}RT} by mol, so again, the internal energy will depend only on temperature and not on volume or pressure.
Distinction between temperature, heat, macroscopic kinetic energy and internal energy
Kinetic theory allows a clear distinction to be made between temperature, heat, and internal energy. Temperature (in kelvins) is a measure of the average kinetic energy of individual molecules. Internal energy refers to the total energy of the molecules inside the object. Thus two hot iron ingots of equal mass can have the same temperature; however, two ingots have twice the internal energy of one. Heat refers to a transfer of energy from one object to another as a result of a difference in temperature. The direction of heat flow between two objects depends on their temperatures, not on how much energy each has. Thus, if 50 g of water at 30 °C is brought into contact with (or mixed with) 200 g of water at 25 °C, heat flows from water at 30 °C to water at 25 °C, even though the internal energy of water at 25 °C is much higher since there is more of it. A distinction must also be made between the macroscopic kinetic energy of an object as a whole and the microscopic kinetic energies of its molecules, which constitute the internal energy of the object. The kinetic energy of an object is an organized form of energy related to the orderly motion of molecules in a straight line direction or around an axis. Instead, the kinetic energies of molecules are completely random and highly disorganized.
Transfer of thermal energy
There are three fundamental mechanisms of thermal energy transfer: conduction, convection and radiation.
- Driving is the transmission of energy in the form of heat from one part of the body to another of the same body, or from one body to another that is in physical contact with it, without appreciable displacement of the body particles.
- Convection is the heat transmission from one point to another within a fluid, a gas or a liquid, by mixing one portion of the fluid with another. In natural convection, the movement of the fluid is due entirely to differences of density as a result of temperature differences; in forced convection, the movement is produced by mechanical means. When forced speed is relatively low, it should be understood that free convection factors such as temperature and density differences can have an important influence.
- Radiation is the transmission of thermal energy from one body to another, which is not in contact with it, through the motion of waves through space.
In all heat transfer mechanisms, the rate of cooling of a body is approximately proportional to the difference in temperature that exists between the body and the environment that surrounds it. This fact is known as Newton's law of cooling. In many real situations, all three heat transfer mechanisms occur simultaneously, although some of them may be more dominant than the others.