Molecular Geometry, Polarity, and Valence Bond Theory

Vocabulary practice cards covering electron geometries, molecular polarity, hybrid orbitals, and valence bond theory based on lecture material.

Electron Group

Any lone pair, single bond, double bond, or triple bond that is counted as a single unit of electrons around a central atom.

Five Basic Electron Geometries

The fundamental arrangements of electron groups around a central atom: linear (2 groups), trigonal planar (3 groups), tetrahedral (4 groups), trigonal bipyramidal (5 groups), and octahedral (6 groups).

Equatorial Positions

The positions within a trigonal bipyramidal electron geometry where lone pairs preferentially reside to minimize electron repulsion.

Axial Positions

The two linear positions positioned at 180^\flavors relative to each other along the vertical axis in a trigonal bipyramidal geometry.

Trigonal Pyramidal Molecular Shape

The molecular geometry formed by four electron groups that include three bonding groups and one lone pair, having bond angles near 109.5^\flavors (e.g., in $NH_3$).

Seesaw Molecular Geometry

The molecular shape resulting from five electron groups composed of four bonding groups and one lone pair in an equatorial position.

Nonpolar Bonds (Course Rule)

Bonds defined specifically for the exam as being either between carbon and hydrogen ($C-H$) or between two identical atoms bonded to themselves.

Electronegativity Periodic Trend

The property that increases from left to right across a period and upward toward fluorine on the periodic table.

Nonpolar Molecule with Polar Bonds

A molecule (such as linear $CO_2$) possessing polar bonds whose dipole moments are equal in magnitude and point in opposite directions, canceling each other out completely.

Valence Bond Theory

A bonding model that describes covalent bonds as the overlap of atomic and hybrid orbitals containing localized electron pairs, accounting for real spatial geometries and bond angles.

Hybrid Orbitals

Equivalent orbitals formed by mathematically mixing atomic orbitals ($s$, $p$, and $d$) on a central atom to accommodate its electron groups.

$sp^3$ Hybridization

The set of four equal hybrid orbitals generated by mixing one $s$ orbital and three $p$ orbitals, corresponding to four electron groups and bond angles of approximately 109.5^\flavors.

$sp^3d$ Hybridization

The set of five hybrid orbitals produced by mixing one $s$, three $p$, and one $d$ orbital, used by central atoms with five electron groups in trigonal bipyramidal geometry.

$sp^3d^2$ Hybridization

The set of six hybrid orbitals created by mixing one $s$, three $p$, and two $d$ orbitals, used by central atoms with six electron groups and 90^\flavors bond angles (e.g., $SF_6$).

$sp^2$ Hybridization

The set of three hybrid orbitals formed by combining one $s$ and two $p$ orbitals for three electron groups (trigonal planar), leaving one unhybridized $p$ orbital available for pi bonding.

$sp$ Hybridization

The set of two hybrid orbitals formed by combining one $s$ and one $p$ orbital for two electron groups (linear), leaving two unhybridized $p$ orbitals available for pi bonding.

Sigma ($\sigma$) Bond

A covalent bond formed by head-to-head orbital overlap directly between two nuclei along the internuclear axis; it is the first and strongest bond present in every single, double, or triple bond.

Pi ($\pi$) Bond

A covalent bond formed by the side-to-side overlap of parallel, unhybridized $p$ orbitals above and below the internuclear axis; double bonds contain one pi bond, while triple bonds contain two pi bonds.

Delocalized Molecular Orbitals

The mechanism within valence bond theory used to represent resonance structures by allowing electron density to be shared equally over multiple adjacent atoms.