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The Joule expansion (a subset of free expansion) is an irreversible process in thermodynamics in which a volume of gas is kept in one side of a thermally isolated container (via a small partition), with the other side of the container being evacuated. The partition between the two parts of the container is then opened, and the gas fills the ...
In thermodynamics, the Joule–Thomson effect (also known as the Joule–Kelvin effect or Kelvin–Joule effect) describes the temperature change of a real gas or liquid (as differentiated from an ideal gas) when it is expanding; typically caused by the pressure loss from flow through a valve or porous plug while keeping it insulated so that no heat is exchanged with the environment.
A number of materials contract on heating within certain temperature ranges; this is usually called negative thermal expansion, rather than "thermal contraction".For example, the coefficient of thermal expansion of water drops to zero as it is cooled to 3.983 °C (39.169 °F) and then becomes negative below this temperature; this means that water has a maximum density at this temperature, and ...
Quantity (common name/s) (Common) symbol/s Defining equation SI unit Dimension Temperature gradient: No standard symbol K⋅m −1: ΘL −1: Thermal conduction rate, thermal current, thermal/heat flux, thermal power transfer
Thermodynamic expansion work. The expansion work of the real gas is different than that of the ideal gas by the quantity (). ...
Here, work is entirely consumed by expansion against the surroundings. Of the total heat applied (709.3 kJ), the work performed (202.7 kJ) is about 28.6% of the supplied heat. Isobaric expansion of a gas pressurized to 2 atmospheres by a 10,333.2 kg mass. Like before, the gas doubles in volume and temperature while remaining at the same pressure.
The adiabatic compression of a gas causes a rise in temperature of the gas. Adiabatic expansion against pressure, or a spring, causes a drop in temperature. In contrast, free expansion is an isothermal process for an ideal gas.
In thermodynamic terms, this is a consequence of the fact that the internal pressure of an ideal gas vanishes. Mayer's relation allows us to deduce the value of C V from the more easily measured (and more commonly tabulated) value of C P : C V = C P − n R . {\displaystyle C_{V}=C_{P}-nR.}