Chemistry: Particles in Solids, Liquids and Gases

Particles in Solids, Liquids, and Gases

• Solids: particles are closely packed, tightly bound, and vibrate constantly, with fixed volume and shape.
• Liquids: particles are close together, can slip by each other, and vibrate faster when heated, with no fixed shape.
• Gas particles have high kinetic energy, spreading out and occupying a larger volume, causing gases to expand and fill containers.

Diffusion

• Diffusion is the spontaneous spreading of a substance due to particle movement.
• Diffusion in gases is faster than in liquids.
• Example: ammonia and hydrogen chloride react to form a white cloud of ammonium chloride.

Gas Laws

Boyle's Law

• At constant temperature, the volume of a given mass of gas is inversely proportional to the pressure.
• P x V = constant.

Charles' Law

• At constant pressure, the volume of a given mass of gas is directly proportional to the Kelvin temperature.
• V / T = constant.

Gay-Lussac's Law

• When gases react, the volumes consumed and produced bear a simple whole number ratio.

Avogadro's Law

• Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
• 1 mole of any substance contains 6 x 10²³ particles.

The Mole

• A mole of any substance is the amount containing 6 x 10²³ particles.
• Avogadro's Constant (L) = 6 x 10²³ mol⁻¹.

Molar Volume of Gases

• Standard temperature and pressure (s.t.p.): 1 mole = 22.4 L or 22,400 cm³ or 2.24 x 10⁻² m³.
• Room temperature and pressure: 1 mole = 24 L or 24,000 cm³.

Relative Molecular Mass

• The relative molecular mass of a substance is the average mass of a molecule relative to one-twelfth of the mass of an atom of C-12.
• Can be calculated by adding the relative atomic masses of all atoms in the molecule.
• Examples: copper = 63.5, zinc = 65, iron = 56.

Relative Molecular Mass & Molar Mass

• Molar mass is the mass in grams of one mole of the substance, measured in g mol⁻¹.
• Molar mass = R.M.M. x g mol⁻¹.
• Examples: H₂SO₄ = 98 g mol⁻¹, C₁₂H₂₂O₁₁ = 342 g mol⁻¹.

The Combined Gas Law

• Combines Boyle's Law, Charles' Law, and Avogadro's Law: P₁V₁ / T₁ = P₂V₂ / T₂.

The Kinetic Theory of Gases

• Assumes:
  • Gases are made up of particles with negligible diameters compared to the distance between them.
  • No attractive or repulsive forces exist between particles.
  • Particles are in constant rapid random motion, colliding with each other and the container walls.
  • Average kinetic energy is proportional to the Kelvin temperature.
  • All collisions are perfectly elastic.
• Ideal gas: perfectly obeys all gas laws under all conditions.
• Real gas: deviates from ideal behavior under low temperature and high pressure due to:
  • Non-negligible particle diameters.
  • Attractive and repulsive forces between particles.
  • Imperfectly elastic collisions.

The Equation of State for an Ideal Gas

• PV = nRT, where P = pressure, V = volume, n = number of moles, R = universal gas constant, and T = temperature.

Experiment: Estimation of the R.M.M. of a Volatile Liquid

• Find the number of moles using PV = nRT.
• Calculate the relative molecular mass, Mr = M / n, where M is the mass in grams and n is the number of moles.

This quiz covers the properties of particles in different states of matter, including solids, liquids, and gases. Learn about the behavior of particles and how they change as they transition between states.