Thermal transfer

A thermal transfer, more commonly called heat, is, along with work, one of the modes of internal energy exchange between two systems: it is a transfer of thermal energy that takes place outside thermodynamic equilibrium. There are three types of heat transfer, which can coexist:

• conduction, due to the progressive diffusion of thermal agitation in matter;
• convection, thermal transfer which accompanies the macroscopic movements of matter;
• radiation, which corresponds to the propagation of photons.

The quantity of heat Q is the quantity of energy exchanged by these three types of transfers, it is expressed in joules (J). By convention, Q > 0 if the system receives energy. Thermodynamics is based on the concept of heat to erect the first and second principles of thermodynamics.

The meaning of the word “heat” in everyday language is often ambiguous and confusing, especially with temperature. While it is true that spontaneous heat transfers take place from regions of higher temperature to regions of lower temperature, it is nevertheless possible to achieve thermal transfer from a cold body to the hot body, using a thermal machine such as a refrigerator. Moreover, during a change of state, a pure substance does not change its temperature while it exchanges energy in the form of heat.

The simplest example of a situation involving heat transfer is that of two bodies in contact with different temperatures. The hottest body transfers energy to the colder body by conduction; its temperature decreases, disorder, thermal agitation, decreases. In return, the temperature of the cold body increases, thermal agitation increases within it.

History and evolution of terminology

Heat, in everyday language, is often confused with the notion of temperature. Although very different from a scientific point of view, the two notions are all the same linked to each other and the history of the genesis of thermodynamics has sometimes induced this confusion. Expressions such as “water is hot” could lead to the mistaken belief that heat is a property of the system when it is a transfer of energy (from water, hot, to the surrounding environment, colder). Also, it is incorrect to say “water loses heat” when it cools. The expression “heat transfer” is however a very widespread pleonasm.

Until the 18th century, scientists believed that heat was made up of a fluid called phlogiston (phlogiston theory).

(Schematic flow of energy in a heat engine. )

In the 19th century, heat was assimilated to a fluid: caloric. The progress and success of calorimetry imposed this theory until the middle of the 19th century. This concept is for example taken up by Sadi Carnot: a heat engine can only operate if heat circulates from a body whose temperature is higher to a body whose temperature is lower; reasoning corresponding to an analogy with a hydraulic machine which derives its energy from the passage of water from a reservoir of high altitude to a reservoir of lower altitude.

It is only with the advent of statistical thermodynamics that heat will be defined as a transfer of the thermal agitation of particles to the microscopic level. A system whose particles are statistically more agitated will exhibit a higher equilibrium temperature, defined at the macroscopic scale. Temperature is therefore a macroscopic quantity which is the statistical reflection of the kinetic energies of particles at the microscopic scale. During random shocks, the most agitated particles transmit their kinetic energies to the less agitated particles. The balance of these microscopic kinetic energy transfers corresponds to the heat exchanged between systems made up of particles whose average thermal agitation is different.

The temperature is an intensive state function used to describe the equilibrium state of a system while heat is a transfer of thermal agitation similar to a quantity of energy, associated with the evolution of a system between two distinct or identical states if the transformation is cyclic.