Electronic Theory in Organic Chemistry Flashcards

This set of vocabulary flashcards covers the fundamental concepts of electronic theory in organic chemistry, including quantum numbers, electronic principles, orbital hybridization, and bonding types.

Quantum Number

A number assigned to each shell around the nucleus to indicate its intrinsic nature and the maximum number of electrons it can contain.

Principal Quantum Number ($n$)

Indicates the main energy level (shell) and the size of the orbital, taking whole number values such as $1, 2, 3, 4$ which correspond to the K, L, M, and N shells.

Azimuthal Quantum Number ($l$)

Shows the number of energy sublevels (s, p, d, f) in each electron shell and describes the shape of the orbital, with values ranging from $0$ to $(n-1)$.

Magnetic Quantum Number ($m$)

Describes the orientation of the orbital in space, having integral values ranging from $-l$ through $0$ to $+l$.

Spin Quantum Number ($s$)

Describes the intrinsic spin of an electron about its axis, with values of $-1/2$ for anti-clockwise and $+1/2$ for clockwise spinning.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same values for all four quantum numbers, or that each orbital holds a maximum of two electrons with opposite spins.

Aufbau Principle

The rule stating that electrons occupy the lowest energy orbitals first in the ground state before filling higher energy levels.

Hund’s Rule

States that electrons fill degenerate orbitals (orbitals with equivalent energy) singly before pairing occurs.

Atomic Orbital

The volume of space where there is a probability of finding a particular electron; each orbital can accommodate no more than two electrons.

Hybridization

The mixing or blending of atomic orbitals to form new hybrid orbitals which are symmetrically disposed in space and possess specific energy and shapes.

Catenation

The unique ability of carbon atoms to form bonds with other carbon atoms, resulting in the formation of large classes of organic compounds.

Sigma Bond ($\text{σ}$)

The strongest type of covalent bond formed by the head-to-head overlap of atomic or hybridized orbitals.

Pi Bond ($\text{π}$)

A weaker covalent bond resulting from the lateral or side-to-side overlap of unhybridized p-orbitals.

$sp^3$ Hybridization

The mixing of one s-orbital and three p-orbitals to form four equivalent orbitals in a tetrahedral geometry with a bond angle of $109.5^{\circ}$ and bond length of $154\,nm$.

$sp^2$ Hybridization

The mixing of one s-orbital and two p-orbitals to form three equivalent orbitals in a trigonal planar geometry with a bond angle of $120^{\circ}$ and bond length of $134\,nm$.

$sp$ Hybridization

The mixing of one s-orbital and one p-orbital to form two equivalent orbitals in a linear geometry with a bond angle of $180^{\circ}$ and bond enthalpy of $837\,kJ\,mol^{-1}$.