Chapter 1-7 Review: Kinetic Molecular Theory and Intermolecular Forces (Video)

Vocabulary flashcards covering key concepts from kinetic molecular theory, diffusion/effusion, molar mass, and intermolecular forces as discussed in the lecture notes.

Kinetic Molecular Theory (KMT)

A model of gas behavior stating that gas particles are small, in constant random motion; they have negligible volume, undergo elastic collisions, and the average kinetic energy is proportional to temperature.

Temperature effect on molecular speed

Increasing temperature raises the average speed of gas molecules; lowering temperature slows them down.

Root-mean-square speed (urms)

A measure of the typical molecular speed in a gas; urms = sqrt(3RT/M) for an ideal gas, where M is molar mass.

Kinetic energy formula for a gas

KE = 1/2 m v^2 (or per mole, KE = 1/2 M u^2); relates mass and speed to the molecule’s energy.

Average kinetic energy and temperature

The average kinetic energy of gas molecules is proportional to temperature (for a mole of gas, ⟨KE⟩ ∝ T, e.g., ⟨KE⟩ = (3/2)RT).

Molar mass units (kg/mol vs g/mol)

Molar mass is the mass per mole; for some equations you must use kilograms per mole (kg/mol) rather than grams per mole (g/mol).

URMS derivation concept

URMS relates to temperature and molar mass via KE and distribution of molecular speeds in a gas.

Diffusion

Spontaneous mixing of gas molecules due to random motion when two gases are allowed to mix in the same space.

Effusion

Gas escape through a small hole into a vacuum; rate depends on molecular speed and is slower for heavier gases.

Graham’s law of effusion

Rate1/Rate2 = sqrt(M2/M1); lighter (smaller M) molecules effuse faster; rate is inversely related to the square root of molar mass.

Intermolecular forces (IMF)

Forces that hold molecules together in liquids and solids, including dipole–dipole interactions, hydrogen bonding, and London dispersion forces.

Dipole-dipole interactions

Electrostatic attractions between polar molecules with partial positive and partial negative ends.

Hydrogen bonding

A strong dipole-dipole interaction where H is bonded to N, O, or F and can interact with lone pairs on neighboring molecules; significant in water.

London dispersion forces (Van der Waals forces)

Weak intermolecular forces arising from temporary instantaneous dipoles; present in all molecules and increase with molar mass.

Induced dipole–induced dipole interaction

Another term for London dispersion forces; temporary dipoles induce neighboring dipoles.

Polar vs nonpolar molecules

Polar molecules have permanent dipoles due to uneven electron distribution; nonpolar molecules have no significant permanent dipole.

Electronegativity

Tendency of an atom to attract electrons in a bond; fluorine is the most electronegative in the notes; hydrogen is cited as least in the teaching example.

Most electronegative element (as stated in notes)

Fluorine.

Bond dipole moment

Dipole moment created by unequal sharing of electrons in a bond; points toward the more electronegative atom.

CH4 (methane) geometry and polarity

Methane has a tetrahedral geometry around carbon; overall nonpolar despite C–H bonds due to symmetry.

H2O geometry and polarity

Water is bent (V-shaped) with ~104.5° bond angle; highly polar and capable of extensive hydrogen bonding.

Bond angles comparison

CH4: ~109.5° (tetrahedral); H2O: ~104.5° due to lone pairs compressing the angle.

Hydrogen bonding significance in water

Hydrogen bonding contributes to water’s high surface tension and unique properties; strong intermolecular interactions involving O–H bonds.

Molar mass and molecular speed intuition

At the same temperature, lighter molecules move faster than heavier ones; diffusion/effusion rates reflect this mass dependence.

Temperature, entropy, and speed distribution

As temperature rises, speed distributions broaden and peak shifts to higher speeds due to greater molecular disorder (entropy).

Practical diffusion/effusion problem approach

When solving Graham’s-law-type problems, compare rates (or times) to find the unknown molar mass; lighter gases yield faster diffusion/effusion.