AP Chemistry Unit 3: Physics and Chemistry of Gases

Ideal Gas

A theoretical gas that meets specific criteria of negligible volume, no intermolecular forces, and elastic collisions.

Ideal Gas Equation

The formula PV = nRT which describes the relationship between pressure, volume, temperature, and amount of gas.

Boyle's Law

States that pressure and volume of a gas have an inverse relationship when temperature is held constant.

Charles's Law

States that the volume of a gas is directly proportional to its temperature at constant pressure.

Avogadro's Law

States that equal volumes of gases, at the same temperature and pressure, contain an equal number of molecules.

Dalton's Law of Partial Pressures

States that the total pressure exerted by a mixture of non-reacting gases is equal to the sum of the partial pressures of each gas.

Mole Fraction

The ratio of the number of moles of a component to the total number of moles in the mixture.

Kinetic Molecular Theory (KMT)

A theory that describes the behavior of gases in terms of particles in constant motion.

Elastic Collisions

Collisions between gas particles where there is no net loss of total kinetic energy.

Root Mean Square Speed (u_{rms})

The measure of the speed of particles in a gas, calculated using u_{rms} = sqrt(3RT/M).

Maxwell-Boltzmann Distribution

A statistical distribution of speeds of particles in a gas at a specific temperature.

High Pressure Effect

Under high pressure, gas particles are forced closer, making their volume significant compared to the container.

Low Temperature Effect

At low temperatures, gas particles have lower kinetic energy and cannot overcome intermolecular forces.

Van der Waals Equation

An equation correcting the Ideal Gas Law for real gas behavior, accounting for intermolecular forces and particle volume.

Correction Term a in Van der Waals Equation

A term that adjusts pressure for the effects of intermolecular forces.

Correction Term b in Van der Waals Equation

A term that adjusts volume for the actual size of gas particles.

Kinetic Energy (KE)

The energy an object has due to its motion, proportional to temperature for gases.

Temperature and KE Relationship

The average kinetic energy of gas particles is directly proportional to the Kelvin temperature.

Pressure Units

Different units such as atm, torr, mmHg, or Pa are used in the Ideal Gas Law; the unit chosen affects the value of R.

Partial Pressure Formula

PA = XA * Ptotal, where PA is the partial pressure, XA is the mole fraction, and Ptotal is the total pressure.

Negligible Volume in KMT

The volume of individual gas particles is negligible compared to the total volume of the gas.

Zero Energy in Kelvin

0 K represents absolute zero, while 0 °C is above absolute zero; temperatures must be in Kelvin for gas equations.

Common Mistake: Celsius vs. Kelvin

Using Celsius directly in gas equations instead of converting to Kelvin.

Common Mistake: KE and Speed Confusion

Mistaking the kinetic energy being directly related to mass; all gases have same KE at same temperature.

Common Mistake: Maxwell-Boltzmann Peak Explanation

Assuming the area under the curve increases with peak height; it represents total particle number and must stay constant.

Real vs Ideal Gas Behavior

Real gases deviate from ideal behavior under conditions of high pressures and low temperatures.