A thermodynamic system contains microscopic kinetic and interaction energy collectively described as internal energy $U$.

Heat is energy transferred across a system boundary because of a temperature difference. Work is energy transferred by other macroscopic interactions, such as expansion against a pressure, electrical driving or shaft motion. Heat and work describe transfers, not substances stored inside the system.

First law

Energy conservation for a thermodynamic system can be written

$$\Delta U=Q-W,$$

where $Q$ is heat transferred into the system and $W$ is work done by the system on its surroundings under this sign convention.

Different conventions are used, so the physical meaning of each sign should always be stated.

Heat and work are processes

A system has a state with an internal energy, but it does not contain a particular amount of “heat” or “work”. Two different processes can take the same system between the same initial and final states with different values of $Q$ and $W$, while producing the same $\Delta U$.

Modes of heat transfer

Heat can cross a boundary by conduction, convection or thermal radiation. These mechanisms differ physically, but all represent energy transfer driven by nonequilibrium conditions.

The first law is an energy-accounting rule: it constrains the change in internal energy regardless of which particular transfer mechanisms produced that change.