Organic Chemistry - Structure and Bonding Vocabulary

Essential vocabulary flashcards covering chemical structure, bonding theories, molecular geometry, hybridization, and resonance concepts in organic chemistry.

Tetravalent

Describing an element, such as carbon, that generally forms four covalent bonds.

Trivalent

Describing an element, such as nitrogen, that generally forms three covalent bonds.

Divalent

Describing an element, such as oxygen, that generally forms two covalent bonds.

Monovalent

Describing an element, such as hydrogen or halogens ($X = \text{F, Cl, Br, or I}$), that generally forms one covalent bond.

Isomers

Compounds that have the exact same molecular formula but different structural arrangements of atoms.

Electronegativity

A measure of an atom's intrinsic ability to attract shared electrons in a bond, which increases moving across a period to the right and up a group on the periodic table.

Formal Charge

The electric charge assigned to an atom in a molecule, calculated assuming shared electrons are split equally: $\text{Formal Charge} = \text{valence } e^- - \text{unshared } e^- - \frac{1}{2}(\text{shared } e^-)$.

Wave Function ($\psi$)

A mathematical solution to wave equations in quantum mechanics; the 3D plot of its square ($\psi^2$) defines an atomic orbital.

Electron Density

The probability of finding an electron within a specified region of space around a nucleus; an atomic orbital shape encompasses $90\text{--}95\%$ of this region.

Node

A location or plane in an orbital where the wave function value is zero ($\psi = 0$) and electron density is completely absent.

Valence Bond Theory

A theory of bonding stating that a covalent bond forms when atomic orbitals overlap, with electrons sharing space like constructive overlapping waves.

$\sigma$ Bond (Sigma Bond)

A covalent bond formed by the direct, head-on overlap of atomic orbitals along the internuclear axis, featuring a circular cross-section.

$\pi$ Bond (Pi Bond)

A covalent bond formed by the side-by-side overlap of unhybridized, parallel p orbitals positioned above and below the internuclear plane.

Induction

The withdrawal or shifting of electron density through orbitals toward a more electronegative atom, resulting in a polarized bond dipole.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; states that electron groups around a central atom adopt a spatial geometry that keeps them as far apart as possible to minimize repulsion.

Steric Number

The total count of bonded atoms plus lone pairs attached to a central atom, used to determine hybridization and electron geometry.

Carbocation

A positively charged carbon atom that possesses three bonds, zero lone pairs, and an empty p orbital.

Carbanion

A negatively charged carbon atom that possesses three bonds and one nonbonding lone pair of electrons.

Resonance

A stabilizing structural effect where $\pi$ electrons or lone pairs are spread across adjacent p orbitals, requiring multiple Lewis structures to accurately describe the electron distribution.

Delocalized Electrons

Electrons that are spread across three or more atoms via overlapping p orbitals, providing additional stability to a molecule or ion.

Localized Electrons

Electrons that are restricted to a single atom (as a lone pair) or shared exclusively between two specific bonded atoms.