Learning Outcomes

• State assumptions of the kinetic theory for ideal gases.
• Explain conditions for ideal gas behavior and limitations at high pressures/low temperatures.
• Use the general gas equation $pV = nRT$ in calculations.

Gas Laws

• Physical condition of gases defined by: Pressure (p), Temperature (T), Volume (V), Amount of Particles (n).
• Pressure (p) is force per unit area; increases with number of molecules (n) and temperature (T).

Boyle's Law

• Volume (V) inversely proportional to pressure (p) at constant temperature: $pV = ext{constant}$.
• Experiments show: $p_1V_1 = p_2V_2$.

Charles' Law

• Volume (V) directly proportional to absolute temperature (T) at constant pressure: $V ext{ a } T; ext{ } V/T = ext{constant}$.

Avogadro's Law

• Volume (V) directly proportional to amount of gas (in moles) at constant pressure and temperature.
• Volume ratios in balanced equations reflect mole ratios.

Ideal Gas Equation

• Derived from combining gas laws: $pV = nRT$.
• Important conversions: 1 atm = 760 mmHg = 101 kPa; 1 dm³ = 1000 cm³.

Gas Mixtures and Partial Pressures

• Dalton’s Law: Total pressure $P_T = P_A + P_B + …$.
• Partial pressure from mole fraction.

Deviations from Ideality

• Kinetic theory basis: gas particles do not exert forces, negligible volume, elastic collisions.
• Real gases deviate due to attractive forces and significant particle volume.

Conditions for Ideal Behavior

• High temperature: minimizes intermolecular forces.
• Low pressure: maximizes distance between particles, reducing attraction influence.