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jason.bourne
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why is
du = δq - δw valid for all processes while
du = δq - p dv valid only for quasi-static processes ?
du = δq - δw valid for all processes while
du = δq - p dv valid only for quasi-static processes ?
jason.bourne said:why is
du = δq - δw valid for all processes while
du = δq - p dv valid only for quasi-static processes ?
U is a state function so you don't need to know dW and dQ to determine dU. You just need to know the initial and final states.jason.bourne said:m little bit confused. i hope you understand my question.
du = δq - δw
is valid also for irreversible process.
is it true that δq and δw values have to be given else we cannot determine du unless the process is reversible so that we can use integral p dv and integral T ds ?
The simple answer is that you can't. There are an infinite number of different irreversible processes that can cause a system to go from one state to another, each using different amounts of work and resulting in different heat flows to the surroundings. ΔU = ΔQ - W is always true, but you need to know either ΔQ to find W or W to find ΔQ if the process is not reversible.i mean, if we are given the values of properties at point 1 and 2 in a reversible process, we can evaluate δq and δw using integral p dv and integral T ds.
so how do we determine δq and δw when the process is irreversible?
This is only correct if you are using the most direct reversible path (eg. an isothermal path, adiabatic path).Andrew Mason said:If the path is reversible you just need to know the initial and final states (P,V,T) to determine the work done and heat flow in moving between those two states.
You can determine the work done and heat flow for a reversible path if you are also given the path or the relationship between P and V during the process. For example, you can determine the change in internal energy and work done by an ideal gas undergoing an adiabatic reversible expansion between two states.But if the path is not reversible, you need to know the initial and final states AND the work done in order to determine the heat flow (or the heat flow in order to determine the work done).
It is not quite as simple as that. In irreversible processes the system and immediate surroundings are not in thermodynamic equilibrium so the states are not precisely defined (eg. P and T). Thermodynamics deals with equilibrium states and transitions from one equilibrium state to another. PV=nRT only applies to an ideal gas in thermodynamic equilibrium, for example.jason.bourne said:work done by/on the system on/by the surroundings is equal to
integral of (external pressure i.e, pressure of surroundings ) * (change in volume).
but if the process is quasi-static we can use internal pressure of gas (pressure of the system) instead of exteral pressure coz pint = pext ± dp, where dp is very very small.
does that mean du = δq - p dv (where p is internal pressure ) valid for reversible process while
du = δq - p dv (where p is external pressure) valid for both reversible and irreversible processes ?
The first law of thermodynamics states that energy cannot be created or destroyed, only transferred or converted from one form to another. This means that the total amount of energy in a closed system remains constant.
A quasi-static process is a thermodynamic process that occurs slowly enough for the system to remain in equilibrium at all times. This means that the system is always close to its internal and external equilibrium states, allowing for accurate measurements and calculations.
The first law of thermodynamics is applied to quasi-static processes by using it to analyze the energy changes within a system as it undergoes a slow and reversible process. This allows for the calculation of work, heat, and internal energy changes in the system.
Examples of quasi-static processes include slowly compressing a gas in a piston-cylinder system, slowly heating a substance in a container, or slowly mixing two substances together. These processes occur slowly and smoothly enough for the system to remain in equilibrium at all times.
The first law of thermodynamics is a statement of the conservation of energy, which is a fundamental principle in physics. It states that the total amount of energy in a closed system remains constant, and energy can only be transferred or converted from one form to another. This means that energy cannot be created or destroyed, only transformed.