General Chemistry: Intermolecular Forces, Gas Laws, Phase Changes, and Thermodynamics

Comprehensive vocabulary flashcards covering intermolecular forces, orbital hybridization, ideal gas laws, kinetic molecular theory, phase changes, phase diagrams, and the first law of thermodynamics.

Intermolecular Force

An attractive electrostatic force occurring between neighboring molecules, atoms, or ions as a result of full or partial opposite charges.

Intramolecular Force

The chemical bonding force (such as a covalent or ionic bond) that holds atoms together within a single molecule, which is much stronger than an intermolecular force.

Dipole-Dipole Interaction

An intermolecular attractive force occurring in polar substances where the electron-rich partial negative end ($\delta^-$) of one molecule attracts the electron-deficient partial positive end ($\delta^+$) of an adjacent molecule, with typical strengths between $5\text{--}25\,\text{kJ/mol}$.

London Dispersion Forces

Intermolecular attractions present in all atoms and molecules resulting from temporary, instantaneous dipoles that arise due to the unequal, fluctuating distribution of electrons, typically ranging from $0.05\text{--}40\,\text{kJ/mol}$.

Hydrogen Bond

An unusually strong dipole-dipole attraction ($10\text{--}40\,\text{kJ/mol}$) formed when a hydrogen atom covalently bonded to a highly electronegative atom ($N$, $O$, or $F$) interacts with a lone pair on another electronegative atom.

Boyle's Law

A gas law formulated by Robert Boyle stating that at constant temperature ($T$) and amount of gas ($n$), volume is inversely proportional to pressure: $P_1V_1 = P_2V_2$.

Charles's Law

A gas law formulated by Jacques Charles stating that at constant pressure ($P$) and amount of gas ($n$), volume is directly proportional to absolute temperature in Kelvin: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Absolute Zero

The theoretical minimum temperature at which the volume of an ideal gas extrapolates to zero, equivalent to $-273.15\,^\circ\text{C}$ or $0.00\,\text{K}$.

Avogadro's Hypothesis

The principle stating that equal volumes of different gases at identical temperatures and pressures contain equal numbers of molecules ($V \propto n$).

$sp^3$ Hybridization in Nitrogen

The hybridization of nitrogen's one $2s$ and three $2p$ valence orbitals without electron promotion to yield four degenerate $sp^3$ hybrid orbitals, accommodating three bonding pairs and one lone pair in a trigonal pyramidal molecular geometry with a bond angle of $< 109.5^\circ$.

$sp^3d$ Hybridization

The mixing of one $s$, three $p$, and one $d$ atomic orbital (such as in phosphorus of $PCl_5$) to produce five degenerate hybrid orbitals oriented in a trigonal bipyramidal geometry, leaving four unhybridized $d$ orbitals.

Bond Polarity Continuum

A spectrum depicting the gradual transition of chemical bonds from nonpolar covalent ($\Delta\text{EN} < 0.5$) to polar covalent ($\Delta\text{EN} = 0.5\text{--}1.9$) to ionic ($\Delta\text{EN} \ge 2.0$) based on electronegativity differences.

Ideal Gas Law

The equation of state relating pressure ($P$), volume ($V$), molar amount ($n$), and absolute temperature ($T$) of an ideal gas: $PV = nRT$.

Universal Gas Constant ($R$)

The fundamental physical constant in the ideal gas law, equal to $0.08206\,\text{L}\,\text{atm}\,\text{mol}^{-1}\,\text{K}^{-1}$, $8.314\,\text{L}\,\text{kPa}\,\text{mol}^{-1}\,\text{K}^{-1}$, or $8.314\,\text{J}\,\text{K}^{-1}\,\text{mol}^{-1}$.

Standard Temperature and Pressure (STP)

A standard reference condition for gases defined as a temperature of $273.15\,\text{K}$ ($0\,^\circ\text{C}$) and a pressure of $1\,\text{atm}$ ($101.325\,\text{kPa}$).

Standard Molar Volume

The volume occupied by exactly one mole of any ideal gas at standard temperature and pressure (STP), equal to approximately $22.4\,\text{L}$.

Kinetic Molecular Theory Postulate 1

Gas molecules are relatively far apart with negligible intermolecular attractions, and the volume of individual particles is tiny compared to the total gas volume, proven by the fact that gases are easily compressible.

Kinetic Molecular Theory Postulate 2

Gas particles are in continuous, random straight-line motion at varying speeds and directions, as proven macroscopically by Brownian motion.

Kinetic Molecular Theory Postulate 3

Collisions between gas molecules or with container walls are completely elastic, conserving total kinetic energy, as proven by the absence of pressure loss in a sealed container over time.

Kinetic Molecular Theory Postulate 4

The average kinetic energy of gas molecules is directly proportional to absolute temperature ($T$) in Kelvin, and molecules of all gases have identical average kinetic energies at the same temperature.

Root-Mean-Square Speed ($u_{\text{rms}}$)

The measure of the speed of particles in a gas that relates directly to temperature and molar mass through the expression $u_{\text{rms}} = \sqrt{\frac{3RT}{M}}$.

Enthalpy of Vaporization ($\Delta H_{\text{vap}}$)

The thermal energy required to convert one mole of a liquid to the gaseous state at constant pressure ($\Delta H_{\text{vap}} = 44.01\,\text{kJ/mol}$ for water).

Enthalpy of Fusion ($\Delta H_{\text{fus}}$)

The amount of heat required to convert one mole of a solid substance to the liquid state at constant pressure ($\Delta H_{\text{fus}} = 6.01\,\text{kJ/mol}$ for ice).

Enthalpy of Sublimation ($\Delta H_{\text{sub}}$)

The enthalpy change when one mole of a solid converts directly into gas, governed by Hess's law as $\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}}$.

Heating Curve

A graphical representation showing the temperature change of a substance as heat is added at a constant rate, characterized by sloped single-phase regions and horizontal plateaus during phase transitions.

Phase Diagram

A pressure-versus-temperature coordinate map delineating the thermodynamic equilibrium regions and boundaries between solid, liquid, and gas phases.

Triple Point

The specific temperature and pressure coordinate on a phase diagram at which solid, liquid, and gas phases coexist in thermodynamic equilibrium ($0.01\,^\circ\text{C}$ and $4.58\,\text{mm}\,\text{Hg}$ for water).

Critical Point

The end point of the liquid-gas coexistence curve on a phase diagram, above which the distinct boundary between liquid and vapor ceases to exist, producing a supercritical fluid.

First Law of Thermodynamics

The principle of energy conservation stating that energy cannot be created or destroyed, only changed from one form to another, expressed as $\Delta U = q + w$.

Internal Energy ($U$)

The total microscopic energy of a system, defined as the sum of all kinetic energies ($KE$) and potential energies ($PE$) of its constituent particles.

Thermodynamic System

The specific part of the universe singled out for observation and thermodynamic analysis, separated from the surroundings by a boundary.

Thermodynamic Surroundings

Everything in the universe exterior to the thermodynamic system with which the system can exchange energy or matter.

Endothermic Process

A thermodynamic change in which heat is absorbed by the system from the surroundings, denoted by a positive sign for heat ($+q$).

Exothermic Process

A thermodynamic change in which heat is released by the system to the surroundings, denoted by a negative sign for heat ($-q$).

Thermodynamic Work ($w$)

Energy transfer resulting from an external force applied across a distance, where work performed on the system is designated as $+w$ and work done by the system is $-w$.