Chapter 1: Drawing Organic Molecules and Bonding Concepts

Vocabulary flashcards covering key terms related to condensed vs. Lewis and skeletal structures, lone pairs and formal charges, resonance, hybridization, bond types, bond length/strength, electronegativity, dipoles, and molecular polarity.

Condensed structure

A representation that shows all atoms in a molecule but omits most bond-line drawings; bonds are implied by adjacency.

Lewis structure

A drawing that shows all atoms, all bonds, lone pairs, and formal charges.

Skeletal structure (bond-line structure)

A zigzag representation where corners/endpoints are carbons; hydrogens bonded to carbons are not drawn; double bonds shown; draw all heteroatoms and the hydrogens directly bonded to them.

Lone pairs

Nonbonding pairs of electrons on atoms; often omitted in skeletal structures; counted when determining formal charges.

Formal charge

A bookkeeping charge assigned to atoms in a Lewis structure to reflect electron distribution.

Heteroatom

An atom other than carbon or hydrogen in an organic molecule.

Resonance structures

Different valid Lewis structures for the same molecule; resonance hybrid is their averaged representation.

Hybridization

The mixing of atomic orbitals to form new, equivalent orbitals that participate in bonding.

sp3 hybridization

Mixing of one s and three p orbitals to form four equivalent sp3 orbitals; tetrahedral arrangement; carbon forms four sigma bonds.

sp2 hybridization

Mixing of one s and two p orbitals to form three sp2 orbitals; trigonal planar geometry; leaves one unhybridized p orbital for pi bonding.

sp hybridization

Mixing of one s and one p orbital to form two sp orbitals; linear geometry; used for bonding in programs like acetylene.

Sigma bond (σ)

Bond formed by end-to-end overlap of orbitals; cylindrically symmetrical; all single bonds are σ bonds.

Pi bond (π)

Bond formed by side-by-side overlap of p orbitals; accompanies a sigma in multiple bonds (double/triple).

Bond length

Distance between nuclei of bonded atoms; shorter for higher bond order (single < double < triple).

Bond strength

Bond energy; generally increases with bond order; triple bonds are strongest.

Electronegativity

An atom’s tendency to attract electrons in a chemical bond.

Bond polarity

Polar bonds arise from unequal sharing of electrons due to differences in electronegativity.

Dipole moment

μ = δ × d; measure of charge separation within a molecule; typically expressed in Debye.

Molecular polarity

Overall polarity determined by geometry and bond dipoles; some molecules (like CO2) have no net dipole due to cancellation.

Tetrahedral geometry

Arrangement of four sp3 orbitals toward the corners of a tetrahedron; bond angles ≈109.5°.

Rotation around a single bond

Rotation around a C–C σ bond is possible/easy (e.g., in ethane).

Rotation around a double bond

Rotation around a C=C bond is restricted due to the pi bond (fixed geometry in ethene).

C=C bond

A double bond consisting of one σ bond and one π bond.

Triple bond

A bond consisting of one σ bond and two π bonds.

s-character and bond strength

Higher s-character in hybrid orbitals (sp > sp2 > sp3) leads to shorter, stronger bonds; sp has 50% s-character, sp2 33%, sp3 25%.

Nonpolar bond

A bond with nearly equal sharing of electrons due to similar electronegativities (e.g., C–C, C–H).

Polar bond

A bond with unequal sharing of electrons due to differing electronegativities (e.g., C–O) resulting in a dipole moment.