Lecture Notes on Gases and States of Matter

The flashcards summarize key concepts from the lecture notes on gases, states of matter, intermolecular forces, and phase changes.

Gas

A phase of matter composed of particles that are moving randomly and very fast in their container.

Pressure

The force exerted per unit area by gas molecules as they strike the surfaces around them.

Elastic Collision

A collision where no exchange of energy occurs between gas particles or surfaces.

Inelastic Collision

A collision where an exchange of energy occurs.

Manometer

An instrument used to measure the pressure of a gas trapped in a container.

Boyle's Law

The volume of a gas is inversely proportional to the pressure when temperature and amount of gas are constant.

Charles's Law

The volume of a gas is directly proportional to the absolute temperature when pressure and amount of gas are constant.

Avogadro's Law

The volume of a gas is directly proportional to the number of gas molecules when pressure and temperature are constant.

Ideal Gas Law

The single law combining simple gas laws, expressed as PV=nRT.

Standard Conditions (STP)

A set of agreed-upon conditions for reporting measurements: 1\text{ atm} pressure and 273\text{ K} (0^{\circ}\text{C}) temperature.

Molar Volume at STP

The volume occupied by 1 mole of gas at STP, which is 22.4\text{ L}.

Dalton's Law

The total pressure of a gas mixture is the sum of the partial pressures.

Partial Pressure

The pressure due to any individual component in a gas mixture.

Mole Fraction

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

Vapor Pressure

The partial pressure of water vapor (or other gas) dependent on temperature.

Kinetic Molecular Theory

A model for the behavior of gases where a gas is a collection of particles in constant motion with negligible attraction and elastic collisions.

Root Mean Square Velocity

A measure of the speed of particles defined by the square root of the average of the squares of particle velocities.

Mean Free Path

The average distance a molecule travels between collisions.

Diffusion

The process of a collection of molecules spreading out from high concentration to low concentration.

Effusion

The process by which a collection of molecules escapes through a small hole into a vacuum.

Graham's Law of Effusion

The ratio of effusion rates of 2 gases is inversely proportional to the square root of their molar masses.

Real Gases

Gases that do not behave like ideal gases at high pressure or low temperature because molecules take up space and have attractions.

Van der Waals Equation

An equation for real gases that modifies the ideal gas law to account for molecular volume and intermolecular attractions.

Crystalline Solids

Solid particles arranged in an orderly geometric pattern (e.g., salt, diamonds).

Amorphous Solids

Solid particles that do not show a regular geometric pattern over a long range (e.g., glass, plastic).

Intermolecular Forces

The attractive forces that exist between all molecules and atoms.

Dispersion Forces (London Forces)

The weakest intermolecular forces caused by temporary dipoles; present in all molecules and atoms.

Dipole-Dipole Attractions

Attractive forces between the permanent dipoles of polar molecules.

Hydrogen Bonding

A very strong dipole-dipole attraction that occurs when Hydrogen is bonded directly to Fluorine, Oxygen, or Nitrogen.

Ion-Dipole Forces

Strong attractions present in mixtures of ionic compounds with polar molecules.

Surface Tension

The energy required to increase the surface area of a liquid by a given amount.

Viscosity

The resistance of a liquid to flow.

Capillary Action

The ability of a liquid to flow up a thin tube against the influence of gravity.

Meniscus

The curving of the liquid surface in a thin tube due to the competition between adhesive and cohesive forces.

Miscible Liquids

Liquids that will always dissolve in each other.

Immiscible Liquids

Liquids that do not mix (e.g., water and pentane).

Vaporization

The process where high-energy molecules at the surface escape the liquid and become a vapor.

Condensation

The process where vapor molecules lose energy and are captured back into the liquid.

Volatile

Describes liquids that evaporate easily.

Nonvolatile

Describes liquids that do not evaporate easily.

Heat of Vaporization

The amount of heat energy required to vaporize 1 mole of liquid.

Dynamic Equilibrium

The state in a closed container where the rate of evaporation equals the rate of condensation.

Boiling Point

The temperature at which the liquid's vapor pressure equals the external pressure.

Normal Boiling Point

The temperature at which the vapor pressure of a liquid equals 1\text{ atm}.

Clausius-Clapeyron Equation

A linear equation relating the natural log of vapor pressure to the inverse of absolute temperature.

Critical Point

The temperature and pressure required to produce a supercritical fluid.

Supercritical Fluid

A state formed at high temperature and pressure where the meniscus between liquid and vapor disappears.

Sublimation

The process where surface molecules of a solid break free and become a gas.

Deposition

The capturing of vapor molecules into a solid.

Melting (Fusion)

The process where a solid turns into a liquid as molecules overcome attractions.

Heat of Fusion

The amount of heat energy required to melt 1 mole of solid.

Phase Diagram

A map describing the different states and state changes that occur at various temperature and pressure conditions.

Triple Point

The temperature/pressure condition where all three states (solid, liquid, gas) exist simultaneously.