CHEM2311 Lecture 2: Bonding, Formal Charge, Resonance, and Line Angle Structures

Flashcards covering key vocabulary terms from CHEM2311 Lecture 2 on Bonding, Formal Charge, Resonance, and Line Angle Structures.

Octet Rule

Bonded atoms are especially stable when their valence shell is completely filled with eight electrons, derived from filling s + p subshells.

Exceptions to the Octet Rule

Atoms like Hydrogen (2 e-), Boron and Aluminum (6 e-), and Phosphorus and Sulfur (>8 e-) that do not strictly follow the octet rule.

Covalent Bond

A chemical bond formed by the sharing of electrons between atoms.

Lewis Structure

A representation of the valence electrons in a molecule, showing bonding electrons as lines and non-bonding electrons as dots.

Bonding Electrons

Electrons in a Lewis structure represented by a line, signifying a shared pair of electrons (2 e-) between two atoms.

Non-bonding Electrons (Lone Pairs)

Electrons in a Lewis structure shown as dots, representing a pair of valence electrons not involved in bonding.

Condensed Structural Formula

A chemical formula that implies the connectivity of atoms in larger molecules, which may also imply branching.

Multiple Bonds

Double or triple bonds used in Lewis structures to satisfy the octet rule for atoms that would otherwise have less than an octet; shared electrons count towards both atoms' electron counts.

Bond Length

The distance between the nuclei of two bonded atoms; decreases as the number of electrons in a covalent bond increases (e.g., Triple bonds are shortest).

Bond Strength

The energy required to break a bond; increases as the number of electrons in a covalent bond increases (e.g., Triple bonds are strongest).

Electronegativity

The ability of an atom to distort bonding electrons towards itself on an arbitrary scale (0.5-4, unitless).

Non-polar Covalent Bond

A bond formed when there is a small or zero difference in electronegativity (ΔEneg < 0.4).

Polar Covalent Bond

A bond formed when there is a significant difference in electronegativity (0.4 ≤ ΔEneg < 1.7), leading to unequal sharing of electrons and charge buildup.

Bond Dipole

The buildup of partial positive (δ+) and negative (δ-) charges on two atoms in a bond due to their differences in electronegativity, represented by a dipole arrow.

Ionic Bond

A bond resulting from attractions between an anion and a cation, where electrons are not shared but transferred, typically forming between atoms with a large electronegativity difference (ΔEneg ≥ 1.7).

Formal Charge

An electron accounting formalism that helps compare the number of valence electrons of a single atom to that of an atom bonded within a molecule, calculated as (# valence e-) – (1/2 bonding e-) – (non-bonding e-).

Resonance Structures

Two or more Lewis structures for a molecule where the connectivity is the same but the placement of some delocalized electrons has changed.

Curved Arrow Notation

A convention used to show the movement of delocalized electrons between resonance structures; arrows must start at a lone pair or bond and indicate electron movement, without moving nuclei.

Resonance Hybrid

A blend of reasonable resonance forms that illustrates the true electron distribution in a molecule, showing delocalized electrons spread over multiple atoms using partial bonds and charges.

Non-equivalent Resonance Structures

Resonance forms that do not contribute equally to the overall resonance hybrid; the more stable contributors typically satisfy the octet rule, minimize charge separation, and place negative charges on more electronegative atoms.

Line-Angle Formula (Skeletal Structure)

A structural shorthand where carbon atoms are implied at the end and vertex of every line, heteroatoms are shown explicitly, hydrogens bonded to carbon are implicit, and hydrogens bonded to heteroatoms are explicit.