The first law of thermodynamics states that energy cannot be created or destroyed, only transformed or transferred.Energy does not appear from nowhere and does not disappear without a trace — it only changes from one form to another. The first law of thermodynamics is like a bank account: the change in your savings equals the sum of deposits minus withdrawals. The same principle in physics: the change in internal energy of a system equals the heat added minus the work done.
The law was formed by the efforts of several scientists in the mid-19th century. Julius Robert von Mayer in 1842 first formulated the principle of conservation of energy based on observations of blood color in the tropics. James Prescott Joule in 1843-1845 experimentally measured the mechanical equivalent of heat, proving that heat is a form of energy. Hermann von Helmholtz in 1847 generalized the law to all forms of energy. Rudolf Clausius in 1850 introduced the concept of internal energy and gave the first rigorous mathematical formulation for heat engines.
The law is the basis of all technical thermodynamics: calculation of engine cycles, refrigerators, chemical reactions. It is applicable to macroscopic systems in a state of thermodynamic equilibrium or close to it. Limits: for relativistic processes, the mass-energy equivalence (E=mc²) must be taken into account; for microscopic systems, quantum fluctuations are important.
The human body is a heat engine: all the energy from food eaten is spent on muscle work, heat release, and stored in fat. You never get more energy than you consumed.
| dU | variable | Infinitesimal change in internal energy — internal energy joule |
| δQ | variable | Infinitesimal heat added to the system — heat joule |
| δW | variable | Infinitesimal work done by the system — work joule |
| d | operator | Exact differential of a state function — exact differential |
| δ | operator | Inexact differential for path-dependent quantities — inexact differential |



