Chapter 7: Gases, Liquids, and Solids

Comprehensive vocabulary flashcards from Chapter 7 covering gas laws, states of matter, and intermolecular forces.

Compressibility

A measure of the change in volume of a sample of matter resulting from a change in pressure.

Thermal expansion

A measure of the change in volume of a sample of matter resulting from a change in temperature.

Kinetic energy

Energy that matter possesses because particles are in motion; the amount depends on temperature ($\text{Higher temperature} = \text{higher amount of kinetic energy}$).

Potential energy

Stored energy that matter possesses as a result of its position, condition, and /or composition, dependent on electrostatic interactions between particles.

Electrostatic interaction

An attraction or repulsion that occurs between charged particles; opposite charges attract while like charges repel.

Kinetic Molecular Theory (KMT)

A set of five statements explaining the physical behavior of matter based on the assumption that any particle is in motion unless at absolute zero.

Elastic collisions

Collisions where the total kinetic energy remains constant and no energy is lost.

Inelastic collisions

Collisions where the energy of motion is lost.

Disruptive force

A force, such as kinetic energy, that causes particles to move away from each other; it is temperature dependent.

Cohesive force

A force, such as potential energy, that holds particles together; it is independent of temperature.

Solid state

The physical state characterized by a large amount of potential energy (cohesive forces) relative to kinetic energy (disruptive forces); particles vibrate about fixed sites.

Liquid state

The physical state characterized by potential energy and kinetic energy of about the same magnitude; particles freely slide over one another but remain touching.

Gaseous state

The physical state characterized by a complete dominance of kinetic energy over potential energy; particles are in constant random motion and widely separated.

Kelvin scale conversion

The temperature scale used in gas laws, calculated as $K = ^{\circ}C + 273$.

Pressure

The force applied per unit area on an object resulting from gas particles colliding with container walls.

Pressure at Sea Level (atm)

$$1\,atm$$

Pressure at Sea Level (mm Hg and Torr)

$$760\,mm\,Hg = 760\,Torr$$

Pressure at Sea Level (psi)

$$14.7\,psi$$

Pressure at Sea Level (Pa)

$$101,325\,Pa$$

Boyle’s law

The volume of a fixed amount of gas is inversely proportional to the pressure if the temperature is constant: $P_1 \times V_1 = P_2 \times V_2$.

Charles’s law

The volume of a fixed amount of gas is directly proportional to its Kelvin temperature if the pressure is kept constant: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Combined gas law

The product of the pressure and volume of a fixed amount of gas is directly proportional to its Kelvin temperature: $\frac{P_1 \times V_1}{T_1} = \frac{P_2 \times V_2}{T_2}$.

Ideal gas law

Describes relationships among four variables for a gaseous substance: $PV = nRT$.

Ideal gas constant (R)

$$0.0821\,\frac{L \cdot atm}{mole \cdot K}$$

Standard temperature and pressure (STP)

Conditions defined as $0^{\circ}C$ ($273\,K$) and $1\,atm$ pressure.

Standard molar volume

The volume occupied by one mole of any gas at STP, which is $22.4\,L$.

Dalton’s law of partial pressures

The total pressure exerted by a gas mixture is the sum of the partial pressures of the individual gases: $P_{Total} = P_A + P_B + P_C + \dots$

Partial pressure

The pressure that a gas in a mixture would exert if it were present alone under the same conditions.

Endothermic change of state

A process where heat energy is absorbed, including evaporation, melting, and sublimation.

Exothermic change of state

A process where heat energy is given off, including condensation, freezing, and deposition.

Evaporation

A surface phenomenon where liquid molecules change to the gas phase by overcoming attractive forces; it causes the temperature of the remaining liquid to decrease.

Vapor

A gas that exists at a temperature and pressure that would normally exist as a liquid or solid.

Physical equilibrium

A state in which two opposing physical processes, such as evaporation and condensation, take place at the same rate.

Vapor pressure

The pressure exerted by a vapor above a liquid when they are in physical equilibrium; increases as temperature increases.

Volatile substance

A substance that readily changes from the liquid to the gas state at room temperature due to high vapor pressure.

Boiling point

The temperature at which the vapor pressure of a liquid equals the external (atmospheric) pressure.

Normal boiling point

The temperature at which a liquid boils specifically at $760\,mm\,Hg$ or $1\,atm$.

Intermolecular forces

Electrostatic attractive forces that act between a molecule and another molecule, determining states of matter and boiling points.

Dipole–dipole interaction

An intermolecular force occurring between polar molecules where the positive region of one is attracted to the negative region of another.

Hydrogen bond

An extra-strong dipole–dipole interaction between a hydrogen atom bonded to $F$, $O$, or $N$ and a lone pair of electrons on another $F$, $O$, or $N$ atom.

London dispersion force

A weak temporary interaction caused by the formation of an instantaneous dipole; the only intermolecular force found in nonpolar molecules.