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In the simplest model, a gas is called ideal when its particles are point-like (no volume) and have no interactions. Real gases behave like ideal gases at low pressure (where the particle volume is neglible compared to the total volume) and high temperature (where condensed phases, i.e. interatomic or intermolecular interactions are disfavored). The size-...


4

A straightforward way to evaluate ideality is to compute the compressibility Z: $$Z=\frac{PV_m}{RT}$$ Z equals 1 for an ideal gas, so deviations from this condition serve as a measure of non-ideality. If you examine a plot of compressibility for a real gas you will in general notice the existence of two regimes: at low pressure the compressibility is ...


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Because it gives simpler-to-derive laws which are often very good approximations Clearly real gases do not always follow the ideal gas laws. They mostly liquefy under some conditions, for example, and, under those conditions they are clearly not ideal. But in practice gas laws are used for things far away from those non-ideal regions. When we are applying ...


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In general for a perfect or ideal gas, $$C_p=C_V + R'$$ (using your notation) where the heat capacities are molar quantities. It follows that for a perfect gas mixture $(C_p)_\text{mix}=(C_V)_\text{mix} + R'$.


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Diffusion is a fundamental molecular phenomenon and it does not have a direction. CASE-1 : I cut open the top of cylinder, now the gas inside experiencees a uniform 1 atm pressure and exerts the same pressure too. As there is no pressure difference so Carbon dioxide should not diffuse out! Why not? Recall if you open a bottle of perfume in a room, ...


2

Sander (Ref. 1) has compiled a useful review of Henry's law constants in water that includes an introduction showing notation and conversions. There are two types of Henry's law constants: Solubility constants convert from pressure to concentration in solution (solubility): $$c=Hp$$ Volatility constants convert from concentration in solution (solubility) ...


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The temperature and the volume of the inner ear are constant. When your ears pop during descent, air from the cabin goes into the ear, increasing the pressure. The law is the following: $$n / P = const$$ You can derive this from the ideal gas law. It has no special name.


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We know from the first Newton motion law, that the net force acting on an object in rest must be zero. The forces acting on the piston are gravity and gas pressure: $$\vec F_g + \vec F_\mathrm{p,down} + \vec F_\mathrm{p,up}=\vec 0 \tag{1}$$ If $V$ is the given bottom gas volume, $V_0$ is the total gas volume, $n$ is the molar amount of each of gases, the gas ...


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When real gases are at high temperature, the kinetic energy prevents any gas particles from interacting via intermolecular forces. With low pressure, the gas particles are separated enough that the intermolecular forces are sparse, therefore, giving rise to the ideal behavior since ideal gases are defined as non-interacting particles. When real gases are at ...


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The van Der Waal equation is $$\left(p+\frac a{V_\mathrm m^2}\right)(V_\mathrm m-b)=RT$$ Here $V_\mathrm m$ is molar volume. When pressure is low and temperature is very high, we can qualitatively say that the molar volume will be very large. Due to this the volume occupied by the molecules (given by $b$) becomes insignificant. The pressure is low but the ...


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I also agree with Karsten Theis that $\ce{He}$ would show more ideal behavior than that of $\ce{H2}$ for longer temperature range, just based on their boiling points. The boiling point of helium (~$\pu{4 K}$ or $\pu{-269 ^\circ C}$ at $\pu{1 atm}$) is more close to absolute zero temperature ($\pu{0 K}$) than that of $\ce{H2}$ (~$\pu{20 K}$ or $\pu{-253 ^\...


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