In free expansion the system is thermally isolated so $\Delta Q = 0$ throughout the process. This implies that process is adiabatic and follows its equation:

$$P_1V_1^\gamma = P_2 V_2^\gamma$$

But since inital and final temperature is same

$$P_1 V_1 = P_2 V_2$$

Both equations can only hold when $V_1 = V_2$. But that is clearly false. How do I resolve this paradox?

  • $\begingroup$ The confusion arises from historic and modern practice. In the past 'adiabatic' meant "without heat transfer". Nowadays, in physics at least, 'adiabatic' almost always is taken to mean "reversible and without heat transfer" which is the same as "at constant entropy". In this modern sense, the free expansion is not adiabatic. $\endgroup$ – Andrew Steane Dec 9 '19 at 16:13

The equation




Is for a reversible adiabatic process. The free expansion is an irreversible process.


Only defines the end points at equilibrium for an ideal gas. It is not the same as


which describes an isothermal process, the path between the end points. The path between the end points is not defined for a free expansion.

Bottom line: A free expansion in an insulated system is adiabatic, but it is not reversible adiabatic.

Hope this helps.


You appear to be confusing the distinctions between what applies at a state versus what applies along a path, as well as between what happens in reversible and irreversible processes.


Free expansion is a process where the external pressure is zero. Adiabatic expansion is a process with no heat flow between the system and its surroundings. Reversible processes keep the system and surroundings in thermal and mechanical equilibrium at all points during the process.

Free expansion is irreversible. It can be done adiabatically. Adiabatic expansion can be done using a reversible or an irreversible process.

A substance at equilibrium has a defined mechanical equation of state. For an ideal gas, it is $pV = nRT$.

During an irreversible process on a closed system containing an ideal gas, we cannot determine the pressure or density for the gas along the path. Along a reversible path, the equation of state of an ideal gas is constrained as

$$pV^n = \mathrm{constant}$$

where $n$ depends on the restrictions isothermal ($n = 1$), isobaric ($n = 0$), isochoric ($n = \infty$), or adiabatic ($n = \gamma$).


Allow an ideal gas to expand with zero external pressure adiabatically.

Free expansion is irreversible. The expression $pV^\gamma = \mathrm{constant}$ does not apply. The mechanical work done on/by the gas is zero in free expansion since $\delta w = - p_{ext} dV = 0$. The heat flow is zero along the adiabatic path. From the first law, $\Delta U = \delta q + \delta w = 0$. For an ideal gas, internal energy change is only a function of temperature

$$\Delta U = \int C_V(T)\ dT $$

Since $\Delta U = 0$, then the gas has the same temperature at the final state as it did at the initial state. This does not mean, the gas has the same, constant temperature at all points along the irreversible path. Indeed, we have no clue what the internal temperature, pressure, and density of the gas are along the irreversible path.


To add to the existing answer:
Ultimately, the relation between initial and final volume/pressure is process-dependent. Adiabatic processes requires zero heat flux, hence if your gas cannot cool into the environment at all, it will change its parameters according to $PV^\gamma=const.$

If the gas can cool instantly into the environment, with the heat flux assumed to be infinite, then the temperature remains the same pre/post compression and your gas parameters evolve according to $PV=const.$

So it is really a question of how the gas can cool away the compressional heat. To put this into a more physical context, in reallife, one will compute the cooling timescale of the gas $t_{\rm cool}$ (e.g. radiative or conductive) and compare it to the timescale for compressing the gas $t_{\rm compression}$. The ratio of those two numbers will decide if one is in either extreme, of zero or infinite cooling, or somewhere in between, in which case one has to put in more work to compute the final state of the gas.


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