Thermodynamics || Lec # 2 || 1st Law of Thermodynamics || Internal Energy || Dr. Rizwana Mustafa

First Law of Thermodynamics is framed as a strict accounting rule for energy: energy cannot be created or destroyed during any process, and it can only change form. That conservation principle matters because it explains why a system’s behavior—like motion, heating, or expansion—must follow a balance between energy transferred as heat and energy transferred as work.

A concrete example makes the conservation idea tangible: a ball thrown upward slows as it rises and speeds up as it falls. Its kinetic energy decreases on the way up while potential energy increases, then the pattern reverses on the way down. Even though the ball’s speed differs between the upward and downward trip, the First Law still demands that energy isn’t disappearing; something else is consuming the lost mechanical energy. In this case, interactions with air particles introduce frictional effects, draining some energy from the ball’s motion and reducing its speed on the return path.

From there, the lecture narrows to thermodynamic systems and internal energy. A thermodynamic system contains particles whose internal energy comes from their microscopic motions and, when relevant, from bonding. For a monatomic gas such as helium, internal energy is tied to translational, vibrational, and rotational contributions (as described in the lecture). For bonded systems—like carbon dioxide or oxygen gas—internal energy also includes bonding energy between molecules, so the system’s internal energy is effectively the sum of all these microscopic energy contributions.

Internal energy change is represented as ΔU (and sometimes ΔE in certain books). When heat Q is supplied to a system, molecular motion increases and internal energy rises. If the system expands against a movable piston, the pressure P acting over a volume change ΔV means mechanical work is done. The lecture uses the relation for work as w = PΔV and then ties the energy flows together: the supplied heat increases internal energy and also provides energy for work. This yields the core balance: ΔU = Q − w. A numerical analogy is used to illustrate the bookkeeping—heat supplied is split into an amount that increases internal energy and an amount spent on doing work.

The First Law is then expressed in alternative forms depending on what happens during the process. If the system does no work, all supplied heat goes into raising internal energy, so ΔE = Q. If no heat transfer occurs and the system is doing work on the surroundings, the internal energy change is driven by work alone, leading to sign conventions where work done by the system is treated as negative in the balance. Finally, the lecture emphasizes that measuring changes in internal energy—via calorimetry and work/heat measurements—lets thermodynamicists predict how a system responds to heating, compression, expansion, or mechanical interaction, all while respecting the conservation rule at the heart of the First Law.