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The ideal-gas law can be employed in general as a reasonable approximation of the fluid thermodynamics for low pressures and high temperatures. Otherwise, intermolecular forces and dimension of fluid particles, which are neglected in the ideal-gas approximation, become relevant and can significantly affect the fluid behavior. [ 4 ]
Real gases are non-ideal gases whose molecules occupy space and have interactions; consequently, they do not adhere to the ideal gas law. To understand the behaviour of real gases, the following must be taken into account: compressibility effects; variable specific heat capacity; van der Waals forces; non-equilibrium thermodynamic effects;
Their predictions are the same for ideal gases. However, for real (non-ideal) gases, the results differ. [3] Dalton's law of partial pressures assumes that the gases in the mixture are non-interacting (with each other) and each gas independently applies its own pressure, the sum of which is the total pressure.
The laws describing the behaviour of gases under fixed pressure, volume, amount of gas, and absolute temperature conditions are called gas laws.The basic gas laws were discovered by the end of the 18th century when scientists found out that relationships between pressure, volume and temperature of a sample of gas could be obtained which would hold to approximation for all gases.
In an ideal gas the particles are assumed to interact only through collisions. The equipartition theorem may also be used to derive the energy and pressure of "non-ideal gases" in which the particles also interact with one another through conservative forces whose potential U(r) depends only on the distance r between the particles. [6]
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Dalton's law (also called Dalton's law of partial pressures) states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases. [1] This empirical law was observed by John Dalton in 1801 and published in 1802. [2] Dalton's law is related to the ideal gas laws.
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