Understanding Real Gas Behavior: Deviations from Ideal Gas Equation

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Continuing to lower the temperature of a real gas, specifically an air mixture of nitrogen and oxygen, leads to significant deviations from ideal gas behavior at temperatures of 90 K and 77 K. These temperatures correspond to the boiling points of nitrogen and oxygen, respectively. As the temperature decreases, gas molecules transition from the vapor phase to the liquid phase, resulting in sudden drops in pressure. This phenomenon highlights the limitations of the ideal gas equation, particularly during phase changes, where the behavior of real gases diverges from theoretical predictions.
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If you had continued to reduce the temperature of your real gas in this experiment (air mixture of mostly nitrogen and oxygen) to lower and lower temperatures, you would observe sudden drops in pressure at 90 K and 77 K. In other words, the behavior of the real gas would deviate significantly from the predicted straight-line behavior of the ideal gas equation determined in Question 3 and your extrapolated graph. Why ? [Hint: Consider what happens to water vapor(H2O(g)) when it is cooled to 0.0 C.]

Could anyone explain this, please ?
 
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Question 3

Where is this mysterious question 3...
 
Phase change

The temperatures 77 & 90 °K correspond to the liquefaction (boiling points) of Oxygen & Nitrogen respectively.

The pressure drops due to the removal of gas molecules from the vapour phase into the more condensed phase of liquid.
 

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