Nature of Heat the Nature of Heat

Nature of Heat

Heat

The nature of heat -- where it comes from, what it is made of, how it moves -- has been a source of fascination to philosophers and scientists since the earliest civilizations. The Ancient Greeks connected heat to their early atomic theory. Natural philosophers and chemists during the Enlightenment in Europe considered heat to be its own substance known as 'caloric.' It was not until the 19th century that physicists connected heat to the emerging theories of energy. In the 1840s, James Joule discovered that the appearance and disappearance of heat was always accompanied by the appearance and disappearance of kinetic energy (Tippler, 1999). It soon was confirmed that heat is not in fact its own substance but is a form of energy.

The study of heat as a form of energy, known as thermodynamics, is closely tied to the kinetic theory of matter. The kinetic theory of matter explains the three states of matter by examining the motion of molecular particles. Thermodynamics also examines the motion of particles to explain the expression of heat in matter.

The heat of an object is associated with the total kinetic energy of the particles within that object, but is not the same thing as the total kinetic energy of the particles within an object. One of the most important aspects of Joule's discovery that work and heat are related was the understanding that heat is not kinetic energy in general but is instead the kinetic energy that is transferred among and between molecules or bodies. According to E. Guha (2000), heat is "energy in transit."

Because heat is a form of energy, it shares the same fundamental properties as energy, including conservation. The principle of the conservation of energy is in fact expressed as a thermodynamic principle known as the First Law of Thermodynamics, which states that energy can change form, but it cannot be created or destroyed. This is true of all energy, and it was the discovery that it was true of heat as well, that heat is never lost but only dissipated, that confirmed heat as a form of energy.

One point of possible confusion is the difference between heat and temperature. While the heat of an object is the amount of transferable kinetic energy among total particles in an object, the temperature of an object is the measurement of the average kinetic energy of molecules within an object with regards to its relationship to a state of thermal equilibrium. A precise definition of temperature is "the property that determines whether or not a body is in thermal equilibrium with its neighboring systems" (Jha, 2004). Because temperature is used as a comparative measure of heat energy, it is important that there be a universal scale of temperature measurement. The measurement of temperature, known as thermometry, with its primary instrument being, of course, the thermometer. When the Kelvin scale of temperature, which is based on the behavior of molecules of an ideal gas, is used, temperature is directly proportional to the average kinetic energy of molecules within the body. In other words, when energy is added to a system, its temperature in Kelvin rises at a directly proportional rate.

While it is a common misconception that heat is measured in temperature, this is not in fact the case. Heat, as has already been stated, is a form of energy and therefore is measured in Joules and/or other energy measurements. This is important because heat can change into other forms of energy (for instance, some heat may translated into the work required to break down molecules in the change from one state of matter to another), and therefore the units used to measure heat must also be able to be used to measure other types of energy. Temperature does not have this flexibility, since it is specifically tied to thermal equilibrium (Zobel, 2010).

The proportional relationship between heat and temperature can be explained by the way heat is related to the kinetic energy of molecules. Within all known substances, molecules are in motion. While the molecules are all traveling at different speeds, the average motion of all molecules in a body is known as the average kinetic energy of the body (Ibid). If energy is added to a system or body, the motion of the molecules increases and the average kinetic energy of the system increases. The amount of heat in a system is the amount of transferable energy, or "work," available in a system. This transferable energy is related to the motion of the molecules in that the motion of the molecules transfers energy among the molecules themselves and to neighboring systems and bodies. The fact that this transfer of energy always occurs in one direction, from workable energy to non-workable energy, is expressed in the Second Law of Thermodynamics, which can be stated in simplistic terms as "heat always dissipates" (Guha, 2000).

The fact that heat always dissipates can be seen in the tendency of bodies at different thermal states to reach a state of thermal equilibrium. If bodies with two different levels of heat are contact with one another, the motion of the molecules within the bodies will transfer energy between themselves until all of the molecules in both bodies and any surrounding systems are at the same average level of kinetic energy -- in other words, until they possess the same level of heat. If the Second Law of Thermodynamics were not true, than the heat in an object would be able to increase or be maintained on its own without any added energy to the system, and the tendency for bodies in contact to reach a state of thermal equilibrium would not be universal.

Different substances require different amounts of energy to reach this equilibrium. This is known as heat capacity. The concept of heat capacity relies on both the measure of heat and the measure of temperature. The heat capacity of an object is the amount of thermal energy required to raise the temperature of an object one degree. It is measured by a tool known as a calorimeter (Jha, 2004).