CH 301 Unit 2 Review: Chemical Bonding, Lewis Structures, VSEPR & Polarity

Vocabulary practice flashcards covering electronegativity, lattice energy, Lewis dot formulas, formal charge calculations, octet exceptions, VSEPR electronic/molecular geometries, and molecular polarity.

Bond Polarity

A property of a covalent bond directly proportional to the difference in electronegativity ($\Delta\text{EN}$) between two atoms and the length of the resulting dipole.

Dipole

A separation of positive and negative electrical charge between two bonded atoms, with a magnitude directly proportional to $\Delta\text{EN}$.

Electronegativity ($\text{EN}$)

A relative scale measuring an atom's ability to attract shared electrons in a bond; row 1 and 2 values include $\text{Li}$ ($1.0$), $\text{Be}$ ($1.5$), $\text{B}$ ($2.0$), $\text{H}$ ($2.1$), $\text{C}$ ($2.5$), $\text{N}$ ($3.0$), $\text{O}$ ($3.5$), and $\text{F}$ ($4.0$).

Lattice Energy ($\text{LE}$)

The energy held within the ionic bonds of a crystalline solid salt, which is directly proportional to the crystal's charge density.

Charge Density

A ratio of the amount of ionic charge (the product of ion charges) to the size of the ions in a crystalline solid salt: $\text{Charge Density} = \frac{\text{amount of charge}}{\text{size of atom}}$.

Sulfate

A polyatomic anion with the formula $SO_4^{2-}$.

Nitrate

A polyatomic anion with the formula $NO_3^-$.

Phosphate

A polyatomic anion with the formula $PO_4^{3-}$.

Carbonate

A polyatomic anion with the formula $CO_3^{2-}$.

Ammonium

A polyatomic cation with the formula $NH_4^+$.

Formula Unit

The lowest whole-number stoichiometric ratio of cations to anions in an ionic compound (e.g., $AB$, $AB_2$, $A_2B_3$) such that total donated electrons equal total accepted electrons.

Formal Charge

The theoretical charge assigned to an individual atom in a Lewis structure, calculated as: $\text{Formal charge} = \text{valence } e^- - [(2 \times \text{nonbonding } e^- \text{ pairs}) + (1 \times \text{bonding } e^- \text{ pairs})]$.

Lone Pairs (Nonbonding Electron Pairs)

Pairs of valence electrons associated with a single atom that do not participate in covalent bond formation.

Bonding Electron Pairs

Pairs of valence electrons shared between atoms to form covalent bonds.

Resonance (Delocalization)

The sharing of electrons across more than two atoms in a molecule via $\pi$ bonds when a single Lewis structure cannot adequately depict the true electron distribution.

$\pi$ (Pi) Bond

A type of covalent bond formed above and below the internuclear axis between atoms that allows electron density to be delocalized across adjacent atoms.

$\sigma$ (Sigma) Bond

A covalent bond formed along the direct internuclear axis between two bonded atoms.

Hypovalent

An octet exception describing a central atom with fewer than 8 valence electrons due to having too few valence electrons to complete an octet, commonly occurring with $\text{Be}$ (forms 2 bonds) and $\text{B}$ (forms 3 bonds).

Radical

A molecule or atom containing an odd total number of valence electrons (such as $NO$ with $11\,e^-$ or $NO_2$ with $17\,e^-$), making it impossible to satisfy the octet rule for every atom.

Expanded Valence

An octet exception where central atoms in period $n \ge 3$ or higher accommodate more than 4 bonds (5 or 6 electron-rich regions, resulting in 10 or 12 valence electrons).

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model predicting the electronic geometry, molecular shape, and bond angles based on minimizing repulsion between electron-rich regions.

Electron-Rich Region

A space around a central atom containing electrons, which can consist of either a single/multiple covalent bond or an unbonded lone pair.

Electronic Geometry

The spatial arrangement of all electron-rich regions (both bonding regions and nonbonding lone pairs) surrounding a central atom.

Molecular Geometry

The three-dimensional spatial arrangement of only the bonded atoms around a central atom.

Linear Geometry

An electronic or molecular shape formed by 2 electron-rich regions with a bond angle of $180^\circ$, or a molecular shape formed by 5 electron-rich regions with 3 lone pairs.

Trigonal Planar

An electronic geometry formed by 3 electron-rich regions with ideal bond angles of $120^\circ$.

Tetrahedral Geometry

An electronic geometry formed by 4 electron-rich regions with ideal bond angles of $109.5^\circ$.

Trigonal Bipyramidal Geometry

An electronic geometry formed by 5 electron-rich regions characterized by bond angles of $90^\circ$, $120^\circ$, and $180^\circ$.

Octahedral Geometry

An electronic geometry formed by 6 electron-rich regions characterized by bond angles of $90^\circ$ and $180^\circ$.

See-saw Geometry

A molecular geometry formed by 5 electron-rich regions consisting of 4 bonded atoms and 1 equatorial lone pair.

T-shaped Geometry

A molecular geometry formed by 5 electron-rich regions consisting of 3 bonded atoms and 2 equatorial lone pairs.

Square Pyramidal Geometry

A molecular geometry formed by 6 electron-rich regions consisting of 5 bonded atoms and 1 lone pair.

Square Planar Geometry

A molecular geometry formed by 6 electron-rich regions consisting of 4 bonded atoms and 2 lone pairs oriented $180^\circ$ apart.

Trigonal Pyramidal Geometry

A molecular geometry formed by 4 electron-rich regions consisting of 3 bonded atoms and 1 lone pair, creating bond angles distorted to less than $109.5^\circ$.

Bent (or Angular) Geometry

A non-linear molecular shape resulting from either 3 electron-rich regions with 1 lone pair or 4 electron-rich regions with 2 lone pairs.

Molecular Polarity

The net dipole moment of a molecule evaluated by three-dimensional symmetry; molecules are nonpolar if symmetrical (possessing three or more planes of symmetry such that $\sum \Delta\text{EN} = 0$) and polar if asymmetrical ($\sum \Delta\text{EN} \neq 0$).