To be specific, heat is the energy transferred between a system and the environment as a consequence of a temperature difference between them. Unlike a mechanical interaction in which work is done, heat requires no macroscopic motion of the system. Instead (we’ll look at the details in Chapter 20), energy is transferred when the faster molecules in the hotter object collide with the slower molecules in the cooler object. On average, these collisions cause the faster molecules to lose energy and the slower molecules to gain energy. The net result is that energy is transferred from the hotter object to the colder object. The process itself, whereby energy is transferred between the system and the environment via atomic-level collisions, is called a thermal interaction.

When you place a pan of water on the stove, heat is the energy transferred from the hotter flame to the cooler water. If you place the water in a freezer, heat is the energy transferred from the warmer water to the colder air in the freezer. A system is in thermal equilibrium with the environment, or two systems are in thermal equilibrium with each other, if there is no temperature difference.

Like work, heat is not a state variable. That is, heat is not a property of the system. Instead, heat is the amount of energy that moves between the system and the environment during a thermal interaction. It would not be meaningful to talk about a “change of heat.” Thus heat appears in the energy equation simply as a value Q, never as $\text{Δ}Q$. FIGURE 19.11 shows how to interpret the sign of $Q$.

FIGURE 19.11 The sign of heat.