Kara Sears
The first law of thermodynamics, formulated in the 19th century, is a statement of the law of conservation of energy in the context of thermodynamic processes. It asserts that energy cannot be created or destroyed in an isolated system, and that the total energy of a system remains constant if no work is done on the system or by the system. This is expressed mathematically as Qsystem + Qsurroundings = 0, indicating that the heat gained by a system equals the heat lost by its surroundings.
| Characteristics | Values |
|---|---|
| Equation | Qsystem = -Qsurroundings |
| Law | First law of thermodynamics |
| Description | The law of conservation of energy when no work is done |
| Description (extended) | The sum of the heat within the system and the heat of the surroundings remains zero |
| Description (mathematical) | qsystem + qsurroundings = 0 |
| Description (general) | The total energy within a system plus the energy in its surroundings remains constant |
| Work done by the system | Absorbed by its surroundings |
| Work done on the system | Absorbed by the system and lost by the surroundings |
| Energy | Cannot be created or destroyed, but can be transformed from one form to another |
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What You'll Learn
- The first law of thermodynamics is a formulation of the law of conservation of energy
- Energy cannot be created or destroyed in an isolated system
- Work done by a system on its surroundings consumes internal energy
- The internal energy of a system is defined by the first law
- The first explicit statement of the first law was made by Rudolf Clausius in 1850
The first law of thermodynamics is a formulation of the law of conservation of energy
The first law of thermodynamics evolved from the discovery that heat and mechanical work are interchangeable forms of energy. This law is expressed in the context of thermodynamic processes, where energy transfer occurs without the transfer of matter. The two principal forms of energy transfer are distinguished as heat and thermodynamic work.
The law also defines the internal energy of a system, which is an important property for understanding the balance of heat transfer, thermodynamic work, and matter transfer into and out of the system. The internal energy of a system can be changed by external factors, such as work done on the system or heat transfer, but the total energy within the system plus the energy in its surroundings remains constant. This is expressed mathematically as Qsystem + Qsurroundings = 0, indicating that energy is conserved.
The first law of thermodynamics is considered the least demanding to grasp conceptually. However, it is a foundational concept in the field of thermodynamics, and it helps to clarify the meaning of energy in this context. It is one of the fundamental laws that govern the behaviour of energy in the universe, alongside the second law of thermodynamics, which states that the sum of the entropies of interacting thermodynamic systems never decreases.
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Energy cannot be created or destroyed in an isolated system
The first law of thermodynamics is a formulation of the law of conservation of energy in the context of thermodynamic processes. This law of conservation of energy states that when no work is done, the heat gained by a system is equal to the heat lost by its surroundings. This can be expressed mathematically as qsystem + qsurroundings = 0. This equation indicates that the total energy within a system plus the energy in its surroundings remains constant, implying that energy is conserved. This is further supported by the fact that in a closed thermodynamic system, if some water is heated, it gains heat from the energy within itself, and the surroundings lose that same amount of energy.
The first law of thermodynamics asserts that energy cannot be created or destroyed in an isolated system. This means that any gain in energy within the system must result in an equivalent loss in the surroundings. This law distinguishes two principal forms of energy transfer: heat and thermodynamic work. It also defines the internal energy of a system, an extensive property that accounts for the balance of heat transfer, thermodynamic work, and matter transfer into and out of the system.
The first explicit statement of the first law of thermodynamics was made by Rudolf Clausius in 1850, referring to cyclic thermodynamic processes and the existence of a function of state of the system, the internal energy. However, notable contributions to the emerging theoretical framework of energy were made as early as the first half of the eighteenth century by French philosopher and mathematician Émilie du Châtelet, who emphasised Leibniz's concept of 'vis viva', mv^2, as distinct from Newton's momentum, mv.
The first law of thermodynamics has important implications for the concept of perpetual motion machines. It states that work done by a system on its surroundings requires the consumption of the system's internal energy, and this loss of internal energy must be resupplied as heat by an external energy source. Therefore, perpetual motion machines of the first kind, which produce more energy than they consume, are impossible according to the first law of thermodynamics.
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Work done by a system on its surroundings consumes internal energy
The first law of thermodynamics, also known as the law of conservation of energy, states that energy cannot be created or destroyed in an isolated system. In other words, the total energy within a system plus the energy in its surroundings remains constant. This principle can be expressed mathematically as qsystem + qsurroundings = 0, where qsystem represents the heat gained by the system and qsurroundings represents the heat lost by the surroundings.
When work is done by a system on its surroundings, it transfers energy to them, leading to a decrease in its internal energy. This can be understood through the equation ΔU = ΔQ - ΔW, where ΔU represents the change in internal energy, ΔQ is the heat put into the system, and ΔW is the work done by the system on its surroundings. If the system performs work, ΔW is positive, indicating an increase in the system's internal energy. Conversely, when work is done on the system, ΔW is negative, signifying a decrease in internal energy.
For instance, consider a closed thermodynamic system where water is heated. The water gains heat from the energy within itself, resulting in a corresponding energy loss in its surroundings. This example illustrates the principle that energy transfers occur between systems without any net creation or destruction of energy.
The relationship between work done by a system and its internal energy can be further examined through the concept of pressure and volume. Assuming constant pressure and volume, any increase in temperature within the system directly contributes to an increase in internal energy. This is because, in a perfect system with no external influences, all the heat gained by the system is derived from its surroundings, resulting in an equivalent loss of heat in the surroundings.
In summary, the first law of thermodynamics emphasizes that the internal energy of a system increases when work is done on it, and decreases when it performs work on its surroundings. This law underscores the fundamental principle of energy conservation, highlighting the interplay between a system's internal energy and its interactions with its surroundings.
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The internal energy of a system is defined by the first law
The first law of thermodynamics, also known as the law of conservation of energy, states that energy cannot be created or destroyed in an isolated system. This is often stated as: "Energy can neither be created nor destroyed in a system of constant mass, although it may be converted from one form to another."
This means that any gain in energy within a system must result in an equivalent loss in its surroundings, and vice versa. In other words, the total energy within a system plus the energy in its surroundings remains constant. This can be expressed mathematically as qsystem + qsurroundings = 0.
Internal energy, a thermodynamic property of a system, is defined by the first law as the sum of the kinetic and potential energies of the particles that form the system. It includes the kinetic energy of molecules and the energy stored in all the chemical bonds between molecules. The internal energy of a system is proportional to its temperature and is a state function, meaning it does not depend on the path by which that state was reached.
Any change in the internal energy of a system is equal to the difference between its initial and final values. This change can be caused by heat transfer or work done by or on the system. The internal energy of a system increases when heat is added and decreases when the system gives off heat or does work. However, since energy is neither created nor destroyed, the change in internal energy always equals zero.
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The first explicit statement of the first law was made by Rudolf Clausius in 1850
The first explicit statement of the first law of thermodynamics was made by Rudolf Clausius in 1850. Clausius was a German scientist and physicist who is best known for his work in thermodynamics, specifically his formulation of the first and second laws of thermodynamics. In 1850, he presented a paper titled "On the Moving Force of Heat and the Laws of Heat which may be Deduced Therefrom," in which he made the first explicit statement of the first law.
The first law of thermodynamics is a formulation of the law of conservation of energy in the context of thermodynamic processes. It is based on the principle that energy cannot be created or destroyed in an isolated system, and that any gain in energy within the system must be accompanied by an equivalent loss in the surroundings. This is expressed mathematically as Qsystem + Qsurroundings = 0, where the heat gained by a system is equal to the heat lost by its surroundings.
Clausius's statement of the first law referred to cyclic thermodynamic processes and the transfer of energy as heat and work. He defined the internal energy of a system, which accounts for the balance of heat transfer, thermodynamic work, and matter transfer into and out of the system. Clausius also introduced the concept of the 'Mean free path' of a particle and made significant contributions to the field of kinetic theory.
While Clausius is credited with the first explicit statement of the first law, it is believed that the law was developed almost simultaneously by other scientists such as Germain Hess, Julius Robert von Mayer, and James Prescott Joule in the 1840s. Additionally, James Watt laid the foundations for the first law in his 1769 patent, which was published in 1774. These contributions helped establish the principles of the first law of thermodynamics, which has become a fundamental concept in the field of thermodynamics and energy conservation.
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Frequently asked questions
The first law of thermodynamics is a formulation of the law of conservation of energy in the context of thermodynamic processes. It distinguishes between two principal forms of energy transfer: heat and thermodynamic work. The law also defines the internal energy of a system, accounting for the balance of heat transfer, thermodynamic work, and matter transfer into and out of the system.
The equation for the first law of thermodynamics when no work is done is expressed as Qsystem + Qsurroundings = 0. This indicates that the heat gained by the system equals the heat lost by its surroundings, resulting in a net energy change of zero.
The equation Qsystem = -Qsurroundings represents the concept that the work performed by a system is absorbed by its surroundings, and vice versa. It illustrates that the heat gained or lost by the system is equal and opposite to that of its surroundings.
The first law of thermodynamics aligns with the principle of energy conservation by stating that energy cannot be created or destroyed in an isolated system. Any gain in energy within the system must be offset by an equivalent loss in the surroundings, emphasizing the conservation of energy.
