MCQ Detailed View

In physical chemistry, the behavior of gases is described by two main models: ideal gas behavior and real gas behavior. The ideal gas law assumes that gas molecules do not interact with each other and occupy no volume. However, real gases deviate from this behavior, especially under extreme conditions of temperature and pressure.

The ideal gas equation is:

PV = nRT

Where:

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  P = Pressure

  
  


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  V = Volume

  
  


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  n = Number of moles

  
  


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  R = Gas constant

  
  


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  T = Temperature

  
  

Real gases approximate ideal behavior best at high temperatures and low pressures. Here’s why:

1. High Temperature:
At high temperatures, gas molecules possess greater kinetic energy. This increased motion overcomes intermolecular attractions (like Van der Waals forces), making the gas behave more ideally.

2. Low Pressure:
At low pressure, gas molecules are farther apart. As a result, the volume of individual molecules and intermolecular forces become negligible. This reduces deviations from ideal behavior.

In contrast:

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  At low temperatures, gas particles move slowly and experience stronger intermolecular attractions, causing deviations from ideal gas law.

  
  


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  At high pressures, gas particles are compressed into a smaller volume, and the finite size of molecules becomes significant.

  
  

Therefore, the most favorable conditions for real gases to behave ideally are high temperature and low pressure.

These concepts are essential when studying gases in physical chemistry and are applied in areas like thermodynamics, gas laws (Boyle’s, Charles’s, and Avogadro’s laws), and chemical engineering processes.