Gas Laws, Kinetic Molecular Theory, Entropy, and Molecular Structure Vocabulary

Vocabulary flashcards covering key terms from kinetic theory, gas laws, spontaneity and entropy, VSEPR and molecular geometry, and intermolecular forces.

Kinetic energy

Energy of gas molecules due to their motion; increases with temperature, leading to higher average speeds and more collisions with container walls at constant volume.

Amontons’s law

At constant volume, pressure is directly proportional to temperature; P1/T1 = P2/T2.

Charles’s law

At constant pressure, volume is directly proportional to temperature; V1/T1 = V2/T2.

Boyle’s law

At constant temperature and amount of gas, pressure times volume is constant; P1V1 = P2V2.

Avogadro’s law

At constant pressure and temperature, volume is proportional to the number of gas molecules; V ∝ n (V = kn).

Dalton’s law

The total pressure of a gas mixture equals the sum of the partial pressures of the individual gases.

Kinetic Molecular Theory (KMT)

A microscopic model explaining gas behavior using postulates about molecular motion and interactions.

Maxwell–Boltzmann distribution

Describes the distribution of speeds of molecules in a gas, showing a range from slow to fast speeds.

Root mean square speed (urms)

The square root of the average of the squared speeds; relates to average kinetic energy and is inversely proportional to the square root of molar mass.

Graham’s law

Effusion rate is inversely proportional to the square root of the molar mass (derived from KMT).

Ideal gas law

Equation PV = nRT relating pressure, volume, temperature, and amount for ideal gases.

Combined gas law

Relates P, V, and T when n is constant or changes are considered: P1V1/T1 = P2V2/T2.

Barometer

Instrument used to measure atmospheric pressure.

Manometer

Instrument used to measure the pressure of a gas in a container; can be closed-end or open-end.

Pressure

Force per unit area; units include Pa, kPa, bar, psi, and atm.

Spontaneity

Spontaneous processes occur naturally under certain conditions without continuous energy input and are not determined by rate.

Reversible process

A process that can proceed in both forward and reverse directions while maintaining equilibrium.

Irreversible process

A process that cannot be reversed by a simple change in conditions.

Entropy (S)

State function measuring disorder or dispersal of matter and energy; increases with more microstates; S solid < S liquid < S gas.

State function

Property that depends only on the current state, not on the path taken to reach that state.

Microstates

Possible quantum or molecular arrangements that correspond to a given macrostate; more microstates mean higher entropy.

VSEPR theory

Predicts the three-dimensional arrangement of electron pairs around a central atom to determine molecular shape.

Electron-pair geometry

Arrangement considering all electron pairs (bonding and lone pairs) around a central atom.

Molecular structure

Placement and connectivity of atoms within a molecule.

Bond dipole moment

Separation of charge within a covalent bond, shown as a dipole along the bond.

Molecular dipole

Net separation of charge in a molecule resulting in overall polarity.

Polar molecule

Molecule with a net dipole moment due to uneven distribution of charge.

Nonpolar molecule

Molecule with no net dipole moment; overall symmetry or equal sharing of electrons yields no polarity.

Intermolecular forces (IMFs)

Attractive forces between molecules that determine physical properties like boiling point and viscosity.

Dispersion forces

Also called London forces; present in all condensed phases; arise from temporary fluctuations in electron distribution and are stronger for larger, more polarizable molecules.

Dipole–dipole attractions

Attractions between permanent dipoles in polar molecules due to partial positive and negative ends.

Hydrogen bonding

A strong type of dipole–dipole interaction occurring when hydrogen is bonded to F, O, or N.