Glossary
Attractive forces (Intermolecular Forces - IMFs)
Forces of attraction between gas particles that become significant at low temperatures, causing real gases to deviate from ideal behavior by reducing collisions with container walls.
Example:
At very low temperatures, the attractive forces between CO2 molecules cause them to condense into dry ice, showing their non-ideal behavior.
Average kinetic energy
The average energy of motion of gas particles, which is directly proportional to the absolute temperature of the gas.
Example:
If you heat a gas, its average kinetic energy increases, causing the particles to move faster and collide more frequently.
Diffusion
The process by which gas particles spread out from an area of higher concentration to an area of lower concentration, resulting in the uniform mixing of gases.
Example:
When you open a bottle of perfume, the scent diffuses throughout the room as the perfume molecules spread out.
Effusion
The process by which gas particles escape through a tiny opening into a vacuum, driven by the pressure difference.
Example:
A helium balloon slowly deflates over time due to the effusion of small helium atoms through tiny pores in the balloon's material.
Elastic collisions
Collisions between gas particles where no net loss of kinetic energy occurs, meaning the total kinetic energy of the system remains constant.
Example:
When billiard balls collide and bounce off each other without losing energy to heat or sound, they are undergoing elastic collisions, similar to ideal gas particles.
Graham's Law of Effusion
A law stating that the rate of effusion of a gas is inversely proportional to the square root of its molar mass, meaning lighter gases effuse faster.
Example:
Using Graham's Law of Effusion, we can predict that hydrogen gas will effuse significantly faster than oxygen gas because hydrogen has a much smaller molar mass.
High pressures (deviation condition)
A condition where real gases deviate from ideal behavior because the volume of the gas particles themselves becomes a significant fraction of the total container volume.
Example:
When air is compressed into a scuba tank at high pressures, the volume of the air molecules themselves can no longer be ignored, leading to non-ideal behavior.
Ideal gases
Hypothetical gases that perfectly follow the assumptions of the Kinetic Molecular Theory, exhibiting predictable behavior under all conditions.
Example:
While no gas is truly ideal, helium at room temperature and low pressure behaves very much like an ideal gas.
Kinetic Molecular Theory (KMT)
A model that describes the behavior of ideal gases based on five fundamental assumptions about gas particles.
Example:
KMT helps explain why a balloon inflates when heated, as the increased kinetic energy of gas particles leads to more frequent and forceful collisions with the balloon's walls.
Larger molecules
Molecules with greater molar mass and typically larger electron clouds, leading to stronger London Dispersion Forces and thus greater deviation from ideal gas behavior.
Example:
Butane (C4H10) is a larger molecule than methane (CH4) and will exhibit more non-ideal behavior because of its stronger London Dispersion Forces.
Low temperatures (deviation condition)
A condition where real gases deviate from ideal behavior because particles slow down, allowing intermolecular attractive forces to become more significant.
Example:
At low temperatures, propane gas in a tank is more likely to behave non-ideally, as the molecules are moving slower and attractions are more pronounced.
Negligible volume (of gas particles)
An assumption of KMT stating that the volume occupied by the gas particles themselves is insignificant compared to the total volume of the container.
Example:
In a large room, the negligible volume of air molecules means we can mostly ignore their individual sizes when calculating the room's total gas volume.
Polar molecules
Molecules with an uneven distribution of electron density, creating partial positive and negative poles, which leads to stronger intermolecular forces and greater deviation from ideal behavior.
Example:
Ammonia (NH3) is a polar molecule and will deviate more from ideal gas behavior than nonpolar methane (CH4) at the same conditions due to its stronger dipole-dipole forces.
Real gases
Actual gases that exist and deviate from ideal behavior, especially under conditions of low temperature and high pressure, due to intermolecular forces and particle volume.
Example:
Water vapor, being a polar molecule, is a real gas that deviates significantly from ideal behavior compared to nonpolar gases like methane.
Van der Waals equation
A modified ideal gas law that includes correction factors for the attractive forces between gas particles and the volume occupied by the particles themselves, accounting for real gas behavior.
Example:
While not used for calculations on the AP exam, the Van der Waals equation conceptually shows how chemists adjust for the non-ideal behavior of real gases.