How is heat (q) defined in thermodynamics?

Heat, denoted as q in thermodynamics, is defined as the transfer of energy that occurs due to a difference in temperature between systems or between a system and its surroundings. This energy transfer can take place in various ways, such as conduction, convection, and radiation, and is fundamentally associated with the movement and interaction of particles.

When two systems at different temperatures come into contact, the energy flows from the hotter system to the cooler one until thermal equilibrium is reached. In this context, heat is not simply the product or outcome of chemical reactions, nor can it be directly equated to the increase in kinetic energy of particles; rather, it is the mode of energy transfer that causes such changes in temperature and state for substances involved.

Additionally, while chemical bonds do store energy, heat itself is not the energy stored in these bonds but the energy that can be transferred as a result of temperature differences, leading to changes in those bonds or the states of the matter in question. Thus, the correct characterization of heat focuses on the process of energy exchange that occurs specifically because of temperature differences.