For something – anything – to acquire a temperature above absolute zero (0 K), it somehow needs to be able to warm. The only real requirement for something to be able to warm is for it to possess a ‘thermal mass’, or simply ‘mass’. A thermal mass provides the thing in question with what is (a bit awkwardly) called a ‘heat capacity’, meaning a capacity to absorb and store energy from some energy source (external or internal).

We already know, from basic thermodynamic principles, how energy can be transferred to (or from) an object. It can be transferred in the form of ‘heat’ [Q] or in the form of ‘work’ [W]. Whenever energy is transferred to an object, the ‘internal energy’ [U] of that object increases as a result, which simply means that the object in question has absorbed (energy isn’t ‘transferred’ to a system until it’s actually become ‘absorbed’ by it) the energy to store it inside its mass, as microscopic kinetic and potential energy of its atoms and molecules.

We already know, from the first post in this series, how system ‘internal energy’ [U] relates to system ‘temperature’ [T]. We know that a system with a high ‘heat capacity’ will warm more slowly than a system with a low ‘heat capacity’, both systems absorbing equal energy inputs, the high-heat-capacity system simply storing a larger portion of the absorbed energy as internal/molecular PE rather than as internal/molecular KE (determining the temperature). Both systems, however, will warm, only at different rates. U and T invariably move in the same direction. Unless there is an ongoing phase transition. Then U will increase and T will not. There is no process, though, where U increases and T decreases. The two correspond.

OK. We know that to make an object warm, we must make it accumulate ‘internal energy’. If it doesn’t, it cannot warm. Continue reading →