0
$\begingroup$

While learning the gaseous state I stumbled upon a question which I don't think has correct answer according to my booklet.

Question:The following equations are occasionally used for approximate calculations on gasses:

Gas(X): $PV_{m}=RT(1+\frac{b}{V_{m}})$
Gas(Y):$P(V_{m}-b)=RT$
Where $V_{m}$ is molar volume, $R$ is universal gas constant, $T$ is temperature, $P$ is pressure, $b$ is van der Waal constant or volume correction term.<
Assuming that gas(X) and (Y) actually obeyed above equation of state then what we can say about comparision of their liquefiabilities.

My answer: According to me neither of the gases would be liquefiable because they have pressure correction term $'a'$ missing so they wouldn't show any intermolecular force and hence their molecules couldn't come together to form liquid.
According to my book it is given that gas(X) is more liquefiable than gas(Y) which seems unusual.
Please tell what is the correct answer and thanks in advance!

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Browse other questions tagged or ask your own question.