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Atom

A basic unit of an element containing positive protons, neutral neutrons, and negative electrons.

Behavior of interacting electrical charges

Opposite charges attract each other, whereas like charges repel each other.

Molecule

A neutral group of two or more atoms joined by covalent bonds, such as $H_2O$ or $O_2$.

Covalent bond

A type of chemical bond formed when atoms share one or more pairs of electrons.

Ion

An atom or group of bonded atoms that carries an overall electrical charge due to the gain or loss of electrons.

Cations vs. Anions

Cations are positively charged ions formed by the loss of electrons; anions are negatively charged ions formed by the gain of electrons.

Electronegativity

A measure of the ability of an atom in a chemical bond to attract shared electrons toward itself.

Lone pair

A pair of valence electrons localized on an atom that is not shared in a covalent bond.

Polar covalent bond

A bond characterized by unequal electron sharing resulting from an electronegativity difference between the bonded atoms.

Significance of $\delta^+$ and $\delta^-$

Symbols representing partial positive and partial negative charges that indicate an uneven electron distribution rather than full ionic charges.

Dipole

A spatial separation of positive and negative electrical charges across a chemical bond or molecule.

Polar molecule vs. Ion

A polar molecule is electrically neutral overall but has an uneven internal charge distribution, whereas an ion carries a net positive or negative charge.

Condition for a molecule with polar bonds to be nonpolar overall

A symmetric three-dimensional geometry that causes individual bond dipoles to cancel out completely, as seen in $CCl_4$.

Reason for the molecular polarity of water ($H_2O$)

Its polar $O-H$ bonds and bent molecular geometry prevent bond dipoles from canceling, creating a net molecular dipole.

Reason for the nonpolarity of oxygen gas ($O_2$)

Both identical oxygen atoms attract the shared electron pair equally due to identical electronegativities, resulting in no dipole.

Procedure for determining molecular polarity

Identify the polar bonds, examine the molecule's three-dimensional shape, and determine whether individual bond dipoles cancel out.

Induced dipole

A temporary uneven distribution of electrons created when a nearby ion or dipole distorts an electron cloud.

Polarizability

The ease with which an electron cloud can be distorted; larger, more diffuse electron clouds are more polarizable.

Intramolecular forces vs. Intermolecular forces

Intramolecular forces exist within a molecule holding its atoms together; intermolecular forces exist between separate molecules.

London dispersion forces

Intermolecular attractions generated when momentary electron fluctuations create temporary dipoles that induce dipoles in neighboring particles.

Molecules exhibiting London dispersion forces

All atoms and molecules, regardless of whether they are polar or nonpolar.

Factors that strengthen London dispersion forces

Greater electron-cloud polarizability and increased surface area contact between adjacent molecules.

Boiling point comparison: straight-chain pentane vs. branched isomers

Straight-chain pentane has a higher boiling point because its elongated shape allows greater surface contact and stronger dispersion forces than compact branched isomers.

Dipole–dipole attractions

Electrostatic attractions between the partially positive pole of one polar molecule and the partially negative pole of an adjacent polar molecule.

Hydrogen bonding

An unusually strong dipole attraction between a hydrogen atom covalently bonded to $N$, $O$, or $F$ and a lone pair on a nearby electronegative atom.

Structural requirement for hydrogen bond formation

The hydrogen atom must be covalently bonded directly to a highly electronegative atom: nitrogen ($N$), oxygen ($O$), or fluorine ($F$).

Self-hydrogen bonding in carbonyl compounds without $O-H$ bonds

Molecules such as methanal possess oxygen but lack $O-H$ bonds, meaning they cannot hydrogen-bond with themselves but can accept hydrogen bonds from water.

Ion–dipole attraction

An electrostatic attraction between a fully charged ion and the oppositely charged partial pole of a polar molecule, such as $Na^+$ and $H_2O$.

Dipole–induced dipole attraction

An attraction occurring when a permanent dipole on a polar molecule distorts the electron cloud of an adjacent nonpolar molecule.

General IMF strength ranking for similarly sized molecules

London dispersion forces < dipole–dipole attractions < hydrogen bonding.

Reason IMF category alone cannot determine attraction strength

Size, polarizability, and molecular shape matter; large nonpolar molecules can have total dispersion attractions that exceed the dipole attractions of smaller polar molecules.

Strategy to identify the IMFs present in a molecular substance

Assign dispersion forces to all molecules, add dipole–dipole attractions if polar, check for hydrogen bonding ($H$ bonded to $N$, $O$, or $F$), and check for ion–dipole attractions in ionic mixtures.

Particle arrangement in solids, liquids, and gases

Solids feature closely packed particles in fixed positions; liquids feature closely packed particles that slide past each other; gases feature widely separated particles in constant, rapid motion.

Microscopic process during evaporation of a molecular liquid

Molecules separate and enter the gas phase by overcoming intermolecular attractions while remaining intact as molecules.

State of covalent bonds during the boiling of water

Intramolecular covalent $O-H$ bonds do not break; only intermolecular attractions between separate $H_2O$ molecules are overcome.

Conservation of molecules during evaporation in a closed container

The total number of molecules remains unchanged assuming no chemical reaction; the molecules simply spread farther apart.

Effect of stronger IMFs on boiling point

Stronger intermolecular forces raise the boiling point because more thermal energy is required to separate the molecules into the gas phase.

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

The quantity of energy required to convert one mole of a liquid into a gas at constant temperature and pressure.

Volatility

The tendency of a substance to evaporate readily, typically reflecting weak intermolecular attractions, low boiling points, and high vapor pressures.

Equilibrium vapor pressure

The pressure exerted by a vapor in dynamic equilibrium with its liquid phase in a closed system, where evaporation and condensation occur at equal rates.

Dynamic nature of vapor-liquid equilibrium

Evaporation does not cease at equilibrium; both evaporation and condensation continue simultaneously at identical rates.

Effect of stronger IMFs on equilibrium vapor pressure

Stronger intermolecular forces lower the equilibrium vapor pressure at a given temperature because liquid particles are held more tightly together.

Importance of temperature when comparing vapor pressures

Heating increases vapor pressure, so substances must be compared at the same temperature to accurately assess differences in intermolecular forces.

Property trends associated with increasing IMF strength

Higher boiling point, higher enthalpy of vaporization ($\Delta H_{\text{vap}}$), lower equilibrium vapor pressure at a given temperature, and lower volatility.

Ionic solid

A repeating three-dimensional lattice of alternating cations and anions held together by strong electrostatic attractions (e.g., $NaCl$).

Molecular solid

A solid composed of discrete molecules held together in a lattice by intermolecular forces, with covalent bonds inside each molecule.

Covalent network solid

An extended three-dimensional structure of atoms connected throughout by strong covalent bonds (e.g., diamond, $SiO_2$).

Metallic solid

A lattice of positive metal cores held together by their attraction to a surrounding sea of delocalized electrons.

Delocalized electrons

Electrons that are not confined to a single atom or bond and are free to move throughout the structure.

Fundamental requirement for electrical conductivity

The presence of mobile charged particles, such as freely moving electrons or ions.

Conductivity comparison: solid $NaCl$ vs. molten/aqueous $NaCl$

Solid $NaCl$ does not conduct because ions are fixed in rigid lattice positions; molten or aqueous $NaCl$ conducts because the ions are mobile.

Electrical conductivity of solid silver ($Ag$)

Solid silver conducts electricity because its delocalized electrons can flow freely throughout the metal lattice.

Electrical conductivity of covalent network solids and the graphite exception

Most network solids do not conduct because electrons are locked in covalent bonds, but graphite conducts because delocalized electrons can migrate across its layers.

Melting point difference: molecular solids vs. covalent network solids

Melting molecular solids breaks weak intermolecular forces, whereas melting covalent network solids requires cleaving strong covalent bonds.

Variation in metal melting points

Melting points of metals vary widely because the strength of metallic bonding depends heavily on electron count and atomic size.

Solute vs. Solvent

The solute is the dissolved substance; the solvent is the dissolving medium (e.g., salt is the solute and water is the solvent).

Solvation process during dissolution

Solute particles separate and become dispersed among solvent particles, forming stabilizing intermolecular attractions with them.

Meaning of the principle 'like dissolves like'

Substances with similar types and strengths of intermolecular forces tend to form solutions with one another; polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.

Orientation of water molecules hydrating dissolved ions

The partially negative oxygen end points toward positive cations, while the partially positive hydrogen ends point toward negative anions.

Solubility of glucose in water

Glucose dissolves readily in water because its multiple $-OH$ groups form extensive hydrogen bonds with water molecules.

Solubility of $NaCl$ in nonpolar $CCl_4$

$NaCl$ dissolves poorly in $CCl_4$ because nonpolar $CCl_4$ cannot form strong ion–dipole attractions to stabilize separated ions.

Solubility comparison: $I_2$ in hexane vs. water

Nonpolar $I_2$ dissolves well in nonpolar hexane via dispersion interactions, but poorly in water because it cannot participate in hydrogen bonding.

Dissolution of nonpolar gases in polar water (e.g., $O_2$)

Small amounts of nonpolar gases can dissolve in water through temporary induced-dipole and dispersion attractions.

Significance of undissolved solid in a solution

Undissolved solid indicates the solution is saturated, not that zero solute has dissolved into solution.

Presence of permanent dipole–dipole attractions: $HCl$ vs. $O_2$

$HCl$ has permanent dipole–dipole attractions because it is polar; $O_2$ is nonpolar and has only dispersion forces.

Self-hydrogen bonding ability: $NH_3$ vs. $CH_4$

$NH_3$ can hydrogen-bond with itself because it contains $N-H$ bonds and a nitrogen lone pair; $CH_4$ has nonpolar $C-H$ bonds and no lone pairs.

Strength of dispersion forces: $F_2$ vs. $I_2$

$I_2$ has stronger dispersion forces because its larger electron cloud is significantly more polarizable than that of $F_2$.

IMF comparison from vapor pressure: $40\,\text{torr}$ vs. $90\,\text{torr}$ at the same temperature

The liquid with a vapor pressure of $40\,\text{torr}$ has stronger intermolecular attractions because fewer molecules escape into the vapor phase.

Solid identity: conducts electricity only after melting

An ionic solid, because melting frees its ions to move and carry an electrical current.

Solid identity: conducts electricity both before and after melting

A metallic solid, because mobile delocalized electrons carry electrical charge in both states.

Melting point contrast: $SiO_2$ vs. $SO_2$

$SiO_2$ is a covalent network solid requiring significant energy to break covalent bonds, while $SO_2$ is a molecular solid held by weaker intermolecular forces.

Common conceptual error regarding melting and boiling molecular substances

Confusing intramolecular covalent bonds with intermolecular attractions; phase changes overcome attractions between molecules while leaving the molecules intact.