States of Matter - Class 11th Chemistry Notes

Introduction

• Lecture by Roshni from LearnoHub
• Overview of the 3 states of matter: Solid, Liquid, Gas
• Focus on detailed concepts, questions, demonstrations, and real-life examples.

Basic Concepts

• States of Matter:
  • Solid: Fixed shape and volume (e.g., a bottle).
  • Liquid: Fixed volume but no fixed shape (takes the shape of the container).
  • Gas: No fixed shape or volume.

Molecules and Their Properties

• Individual molecules have different properties compared to bulk matter:
  • A single molecule of water doesn’t wet the hand, but many water molecules can exhibit wetting properties.
  • Boiling occurs in bulk (e.g., boiling a pot of water) rather than a single molecule.
• Bulk properties depend on collections of molecules.

Intermolecular Forces

• Forces that hold molecules together:
  • Intramolecular Forces: Forces within a single molecule (e.g., between atoms).
  • Intermolecular Forces: Forces between different molecules.
• Types of Intermolecular Forces:
  1. London Forces: Exist due to temporary dipoles in molecules.
     • Important for non-polar molecules.
  2. Dipole-Dipole Forces: Occur between polar molecules (e.g., HCl).
  3. Dipole-Induced Dipole Forces: Occur between a polar and non-polar molecule.

Polar and Non-Polar Molecules

• Polar Molecules: Unequal sharing of electrons leads to partial charges.
  • Example: Water (H2O).
• Non-Polar Molecules: Equal sharing of electrons.
  • Example: Oxygen (O2).

Types of Van der Waals Forces

1. London Dispersion Forces:
   • More pronounced in larger atoms due to larger electron clouds.
   • Acts between non-polar molecules.
2. Dipole-Dipole Forces:
   • Occurs between polar molecules with permanent dipoles.
3. Dipole-Induced Dipole Forces:
   • Interaction between polar and non-polar molecules.

Thermal Energy and Motion

• Thermal Energy: Increases particle motion when heat is added.
• Thermal motion is higher in gases due to greater freedom of movement.

Gaseous State: Properties and Laws

• Properties of Gases:
  • Low density and can be compressed.
  • No fixed shape or volume; take the shape of the container.
  • Exert pressure by collisions with container walls.
  • Mix evenly without mechanical aid.

Important Gas Laws

1. Boyle's Law: For a constant temperature, pressure is inversely proportional to volume.
   • Mathematically: ( P_1 V_1 = P_2 V_2 )
2. Charles's Law: For a constant pressure, volume is directly proportional to temperature.
   • Mathematically: ( V_1 / T_1 = V_2 / T_2 )
3. Gay-Lussac's Law: For constant volume, pressure is directly proportional to temperature.
   • Mathematically: ( P_1 / T_1 = P_2 / T_2 )
4. Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain an equal number of molecules.
   • Mathematically: ( V \propto N )

Ideal Gas Equation

• Ideal gas behavior is defined by: ( PV = nRT ) where:
  • P = Pressure
  • V = Volume
  • n = Number of moles
  • R = Universal gas constant
  • T = Temperature

Real Gases vs. Ideal Gases

• Real Gases exhibit deviations from ideal behavior, especially at high pressure and low temperature, due to intermolecular forces.
• Van der Waals Equation for real gases is used to account for intermolecular forces:
  • Mathematical form involves corrections in pressure and volume.

Compressibility Factor ( Z )

• Measures deviation from ideal behavior. For an ideal gas ( Z = 1 ).
• If ( Z < 1 ), gas is more compressible than expected.
• If ( Z > 1 ), gas is less compressible than expected.

Conclusion

• Summary and encouragement to practice more questions.

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