The First Law of Thermodynamics represents the application of the universal principle of conservation of energy to thermodynamic systems. It establishes a fundamental relationship between heat, work, and internal energy and provides a quantitative description of energy changes during physical and chemical processes.
Before the formulation of the First Law, heat was considered a material substance known as caloric, which was believed to flow from one body to another without being created or destroyed. This view was challenged by experimental observations that showed heat could be produced continuously by mechanical work, such as friction and compression.
These experiments led to the realization that heat and work are not substances but different modes of energy transfer. This understanding required a general law that could unify heat and work under a single principle while preserving the total energy of the system and its surroundings.
Need for the First Law of Thermodynamics
The need for the First Law arose from the failure of the caloric theory to explain the interconversion of mechanical work and heat. Experiments demonstrated that a fixed amount of mechanical work always produced a fixed amount of heat, indicating a direct equivalence between the two forms of energy.
This equivalence made it clear that energy could change its form but could not be created or destroyed. A universal law was therefore required to account for all energy transformations occurring during thermodynamic processes, regardless of the nature of the system.
Statement of the First Law of Thermodynamics
The First Law of Thermodynamics may be stated as follows: energy can neither be created nor destroyed, but it can be converted from one form to another.
In thermodynamic terms, this statement implies that when a system undergoes a change from one state to another, the change in its internal energy is determined by the heat supplied to the system and the work done by the system.
Energy, Heat, and Work
Energy is the capacity to do work or to produce heat and exists in various forms such as mechanical, thermal, chemical, and electrical energy. Thermodynamics does not focus on the detailed nature of these forms but rather on their transformation during processes.
Heat is defined as the energy transferred between a system and its surroundings solely due to a temperature difference. Work is defined as the energy transferred when a force acts through a distance. Both heat and work represent energy in transit and are observed only during a process.
Once heat or work crosses the boundary of a system, it contributes to a change in the internal energy of the system. The system retains no memory of whether the energy entered as heat or left as work.
Sign Convention
To apply the First Law quantitatively, a consistent sign convention is adopted. Heat absorbed by the system from the surroundings is taken as positive, while heat released by the system is taken as negative.
Work done by the system on the surroundings is taken as positive, whereas work done on the system by the surroundings is taken as negative. This convention ensures consistency in the mathematical formulation of the First Law.
Conceptual Form of the First Law
The First Law of Thermodynamics connects internal energy, heat, and work and can be expressed conceptually as follows: the change in internal energy of a system is equal to the heat supplied to the system minus the work done by the system.
This principle applies to all thermodynamic processes, whether they involve physical changes such as expansion and compression or chemical changes such as reactions and phase transitions.
The First Law does not predict the direction of a process or its spontaneity; it merely ensures that energy is conserved in every transformation. These limitations necessitate the introduction of the Second Law of Thermodynamics.