Bonding, Hybridization and Molecular Structure: every key term you need (+ practice quiz)
25 flashcard terms for Organic Chemistry I Topic 1, written to match the course framework. Study them here, then drill them as interactive flashcards, or test yourself with the 15-question quiz — free, no account needed.
An electron bookkeeping value equal to valence electrons minus lone-pair electrons minus half the shared electrons. It identifies which Lewis structure is most reasonable and marks where charge formally sits.
Lewis structure
A drawing that shows every valence electron as a bonding pair or a lone pair. It is the starting point for predicting geometry, polarity, acidity, and where electrons will move in a mechanism.
Octet rule
Second-row elements are most stable holding eight valence electrons. Carbon, nitrogen, oxygen, and fluorine essentially never exceed it, while third-row elements such as sulfur and phosphorus can.
Sigma bond
A bond formed by head-on overlap of orbitals, with electron density concentrated along the internuclear axis. Rotation about a sigma bond does not break the orbital overlap.
Pi bond
A bond formed by side-on overlap of parallel p orbitals, with density above and below the internuclear axis. Rotating about it would break the overlap, so pi bonds lock geometry.
sp3 hybridization
Mixing one s and three p orbitals gives four equivalent orbitals at roughly 109.5 degrees. Typical of saturated carbon, the nitrogen of an amine, and the oxygen of an alcohol.
sp2 hybridization
Mixing one s and two p orbitals gives three orbitals in a plane at roughly 120 degrees, leaving one unhybridized p orbital perpendicular to that plane for pi bonding.
sp hybridization
Mixing one s and one p orbital gives two orbitals 180 degrees apart, leaving two perpendicular p orbitals. Found at the carbons of an alkyne and of a nitrile.
Percent s character
The fraction of an orbital derived from the s orbital: 25 percent for sp3, 33 percent for sp2, 50 percent for sp. Higher s character holds electrons closer to the nucleus and shortens bonds.
Bond dipole
A separation of charge along a single bond caused by unequal electronegativity, drawn as an arrow pointing toward the more electronegative atom. Bond dipoles add as vectors.
Molecular dipole moment
The vector sum of all bond dipoles and lone-pair contributions in a molecule. Symmetric arrangements can cancel completely, which is why carbon tetrachloride is nonpolar.
Electronegativity
The tendency of an atom to pull shared electrons toward itself. It rises going right across a period and falls going down a group, which controls bond polarity and inductive effects.
Resonance structures
Two or more valid Lewis drawings that differ only in the placement of pi electrons or lone pairs, never in the positions of atoms. They are not separate species in equilibrium.
Resonance hybrid
The single real structure that a set of resonance forms approximates, weighted toward the most stable contributors. Delocalization always lowers the energy of the actual molecule.
Major resonance contributor
The drawing with full octets, the fewest formal charges, and any negative charge on the most electronegative atom. It best approximates the true electron distribution.
Curved arrow formalism
A notation where a double-headed arrow shows a pair of electrons moving from a source of density to an electron-poor site. Arrows always start at electrons, never at a positive charge.
Inductive effect
Transmission of charge through sigma bonds by electronegativity differences. It weakens quickly with distance, so a chlorine two carbons away barely affects acidity compared with one bonded directly.
Constitutional isomers
Compounds with identical molecular formulas but different connectivity of atoms. Butane and isobutane are the classic pair, and they differ measurably in boiling point and stability.
Bond-line structure
A shorthand drawing where each vertex and each line end is a carbon and hydrogens on carbon are implied. Heteroatoms and their hydrogens are always written explicitly.
Condensed structural formula
A text representation such as CH3CH2CH2OH that lists atoms in connectivity order without drawing bonds. Parentheses indicate branches hanging off the main chain.
Degree of unsaturation
A count of rings plus pi bonds calculated from the molecular formula. Each degree corresponds to two hydrogens missing relative to the fully saturated acyclic formula.
VSEPR geometry
The prediction that electron groups around a central atom spread as far apart as possible. Lone pairs occupy more angular space than bonding pairs and compress the remaining angles.
Antibonding orbital
A molecular orbital with a node between the nuclei, higher in energy than the atomic orbitals that formed it. Populating it destabilizes and eventually breaks the bond.
Bonding molecular orbital
The lower-energy combination formed when atomic orbitals overlap in phase, concentrating electron density between the nuclei and holding the atoms together.
Bond length and bond order
As bond order rises from single to double to triple, bonds shorten and strengthen. Hybridization contributes too, since higher s character pulls bonded atoms closer.