Chemical Bonding, Molecular Geometry, and Intermolecular Forces

Vocabulary practice flashcards covering electronegativity, bond types, VSEPR theory, molecular shapes, molecular polarity, and intermolecular forces.

Electronegativity

A measure of how strongly an atom attracts shared electrons within a chemical bond. Noble gases typically have no values because they have full outer valence shells.

$\Delta\text{EN}$ Bond Classification Rules

The numerical criteria for determining bond type: $\Delta\text{EN} < 0.5$ indicates a nonpolar covalent bond; $0.5 \le \Delta\text{EN} \le 1.6$ indicates a polar covalent bond; $\Delta\text{EN} > 2.0$ indicates an ionic bond; and between $1.6$ and $2.0$ is ionic if a metal is involved, or polar covalent if only nonmetals are involved.

Nonpolar Covalent Bond

A covalent bond formed between identical atoms or atoms with very similar electronegativities ($\Delta\text{EN} < 0.5$, such as $\text{C}-\text{H}$), resulting in equal sharing of electrons and a symmetrically distributed electron cloud.

Polar Covalent Bond

A covalent bond formed between atoms with differing electronegativities ($\Delta\text{EN}$ between $0.5$ and $1.6$) in which electrons are shared unequally and pulled toward the more electronegative atom.

Dipole

A separation of positive and negative electrical charges across a bond or molecule, represented by an arrow symbol pointing toward the more electronegative atom ($\delta-$) with a crossed tail at the less electronegative atom ($\delta+$).

Partial Charges ($\delta+$ and $\delta-$)

Non-integer charges (less than full ionic charges) that arise in polar covalent bonds: $\delta-$ denotes the atom pulling electrons more strongly, and $\delta+$ denotes the atom with less electron density.

Electron Density

The probability or concentration of finding electrons in a region around an atom or within a bond; it is significantly higher around the more electronegative atom in a polar bond.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory; states that electron domains around a central atom repel each other and orient themselves as far apart as possible in three-dimensional space to minimize repulsive forces.

Electron Domain

A region of high electron density around a central atom, where single, double, or triple bonds and nonbonding lone pairs each count as exactly one domain.

Bonding Electron Domain

An electron domain where a central atom shares valence electrons with a peripheral atom through a single, double, or triple covalent bond.

Nonbonding Electron Domain

An electron domain consisting of an unshared lone pair of electrons localized on the central atom.

Linear Molecular Geometry

A three-dimensional molecular shape characterized by two electron domains on the central atom and a bond angle of $180^\circ$, exemplified by $\text{BeF}_2$ and $\text{CO}_2$.

Trigonal Planar Molecular Geometry

A three-dimensional molecular shape characterized by three bonding electron domains and zero lone pairs on the central atom, with bond angles of $120^\circ$, exemplified by $\text{H}_2\text{CO}$.

Tetrahedral Molecular Geometry

A three-dimensional molecular shape characterized by four bonding electron domains and zero lone pairs on the central atom, with bond angles of $109.5^\circ$, exemplified by $\text{CH}_4$ and $\text{CCl}_4$.

Pyramidal (Trigonal Pyramidal) Molecular Geometry

A three-dimensional molecular shape resulting from four electron domains (three bonding domains and one lone pair) on the central atom, with bond angles compressed to approximately $107^\circ$, exemplified by $\text{NH}_3$ and $\text{PH}_3$.

Bent Molecular Geometry

A three-dimensional molecular shape resulting from either four electron domains with two lone pairs (bond angle approximately $104.5^\circ$, e.g., $\text{H}_2\text{O}$, $\text{H}_2\text{S}$) or three electron domains with one lone pair (bond angle near $120^\circ$, e.g., $\text{SO}_2$).

Line, Wedge, and Dash Drawings

A convention used to represent three-dimensional molecular shapes on flat paper: a solid line represents a bond in the plane of the paper, a solid wedge represents a bond coming forward in front of the plane, and a dashed line represents a bond pointing backward behind the plane.

Molecular Dipole

The overall dipole moment of an entire molecule determined by vector addition of all individual bond dipoles; it dictates the molecule's overall polarity and how strongly it interacts with surrounding molecules.

Polar Molecule

A molecule that has an asymmetric charge distribution, resulting in a distinct partially positive end and partially negative end because its polar bond dipoles do not cancel each other out (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{H}_2\text{CO}$).

Nonpolar Molecule

A molecule that possesses no net molecular dipole because all bonds are nonpolar or because its polar bonds are symmetrically arranged around the central atom, causing the individual bond dipoles to cancel out (e.g., $\text{CO}_2$, $\text{CH}_4$, $\text{BeF}_2$).

Intramolecular Bond

A strong attractive force within a single molecule that holds atoms together (such as a covalent bond); breaking an intramolecular bond changes the chemical identity of the substance and constitutes a chemical change.

Intermolecular Force (IMF)

A relatively weak attractive force that exists between two separate molecules; breaking these forces does not alter the chemical identity of the molecules and constitutes a physical change.

Dispersion Forces (London Forces)

The weakest intermolecular force caused by constant electron motion that temporarily shifts electron density to one side of an atom or molecule, producing a momentary dipole that induces dipoles in neighboring atoms; present in all molecules and noble gases.

Polarizability

The ease with which an atom or molecule's electron cloud can be distorted to create a temporary dipole; larger atoms and molecules with more electrons are more polarizable and exhibit stronger London dispersion forces.

Dipole-Dipole Forces

Medium-strength electrostatic attractive forces occurring between the permanent partially positive ($\delta+$) end of one polar molecule and the partially negative ($\delta-$) end of an adjacent polar molecule.

Hydrogen Bonding

An unusually strong type of dipole-dipole attraction that occurs between a hydrogen atom covalently bonded to a highly electronegative atom ($\text{N}$, $\text{O}$, or $\text{F}$) and an unshared electron pair on an $\text{N}$, $\text{O}$, or $\text{F}$ atom of another molecule.

Ion-Ion Forces

The extremely strong electrostatic attractions between positive cations and negative anions in an ionic compound; stronger than dipole-dipole forces and hydrogen bonding, leading to very high melting and boiling points.

Volatility

The tendency of a substance to vaporize and turn into a gas at room temperature; substances held together by weak intermolecular forces have higher volatility than those held by strong forces.

Covalent Network Solid

A giant single-macromolecule structure formed entirely of continuous covalent bonds ($100\%$ intramolecular attractions and no intermolecular forces), resulting in extreme hardness, insolubility, and extremely high melting points (e.g., diamond melting at $3550\,^\circ\text{C}$, quartz $\text{SiO}_2$).