Introduction to Alkanes and Molecular Orbitals

Vocabulary flashcards covering key concepts from bonding, molecular orbitals, hybridization, and representative hydrocarbons discussed in the lecture.

Molecular orbital

A region in a molecule formed by the combination of atomic orbitals; electrons occupy bonding orbitals (lower energy) or antibonding orbitals (higher energy); each orbital holds up to two electrons.

Sigma bonding orbital

A bonding molecular orbital formed by end-to-end overlap along the bond axis; leads to electron density between the nuclei and can hold up to two electrons.

Antibonding orbital

An antibonding molecular orbital formed by out-of-phase overlap; higher energy than bonding orbitals and contains a node between nuclei; often empty or weakens bonding.

In-phase overlap

Overlap where the wavefunctions align (peaks with peaks); constructive interference that increases electron density between nuclei.

Out-of-phase overlap

Overlap where wavefunctions are opposite (peaks align with troughs); destructive interference that can create a node and weaken bonding.

Phase

The relative position of the peaks and troughs of overlapping waves; matters for interference but does not mean charge itself.

Amplitude

The deviation of a wave from its center line; in chemistry linked to electron density—the larger the amplitude, the greater the density in that region.

Wavelength

The distance between successive peaks of a wave; a property used to describe waves, including electron waves.

Frequency

How many wave cycles occur per unit time; related to the energy of waves (higher frequency means higher energy in photons).

Electron density

A measure of how likely electrons are to be found in a region of space; higher amplitude corresponds to higher electron density.

Node

A point or region where electron density goes to zero due to destructive interference between overlapping orbitals.

Valence electrons

Electrons in the outermost (valence) shell that participate in bonding; chemistry often focuses on these while core electrons stay energetically deep.

Core electrons

Electrons in inner shells that are generally not involved in bonding; they are lower in energy and largely ignored in bonding discussions.

Hybridization

Mixing of atomic orbitals to form new, equivalent hybrid orbitals that explain molecular geometry (e.g., sp3, sp2, sp).

sp3 hybridization

One s and three p orbitals combine to form four equivalent sp3 orbitals; leads to tetrahedral geometry (e.g., CH4) with ~109.5° angles.

sp2 hybridization

One s and two p orbitals combine to form three sp2 orbitals; yields trigonal planar geometry (~120°); leftover p orbital forms a pi bond (as in ethene).

sp hybridization

One s and one p orbital combine to form two sp orbitals; yields linear geometry (e.g., acetylene) with 180° angles and supports a triple bond (one sigma and two pi bonds).

Methane CH4

Carbon forms four sigma bonds using sp3 hybrid orbitals; structure is tetrahedral with an octet around carbon.

Ethene C2H4 (ethylene)

Each carbon is sp2 hybridized; carbons form three sigma bonds in a plane and share a pi bond between them, giving a C=C double bond and ~120° geometry.

Acetylene C2H2

Carbons are sp hybridized; linear molecule with a C≡C triple bond (one sigma and two pi bonds); each carbon also bonds to one hydrogen.

Pi bond

Bond formed by sideways overlap of unhybridized p orbitals; electron density is above and below the bond plane; part of double and triple bonds.

Sigma bond

Bond formed by end-to-end overlap along the bond axis; electron density lies between nuclei and forms the first bond in a pair.

Bond angle 109.5°

Tetrahedral geometry, as in methane, where four substituents around carbon are arranged to give ~109.5° angles.

Bond angle 120°

Trigonal planar geometry (as in sp2 systems like ethene), with bonds arranged at ~120° angles.

Bond angle 180°

Linear geometry (as in acetylene) with substituents arranged in a straight line, 180° apart.

Octet rule

Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell.

Overlap

Spatial coincidence of orbitals from adjacent atoms that leads to bond formation; can produce sigma or pi bonds and their antibonding counterparts.

Octahedral geometry

Six electron domains around a central atom with ~90° angles; discussed as an example and less common in simple organic molecules because d-orbitals are not typically involved.