Chemical Bonding, Lewis Dot Structures, and Periodic Trends

Vocabulary practice flashcards covering electronegativity, bond types, Lewis structure conventions, resonance, formal charges, and octet rule exceptions from the lecture.

Covalent Bond

A type of chemical bond formed between two nonmetals where electrons are shared to minimize energy, achieve stability, and attain a full valence shell of eight electrons.

Ionic Bond

A chemical bond formed between a metal and a nonmetal where electrons are completely transferred, creating positive and negative ions.

Electronegativity

A periodic trend describing the strength or ability of an atom to pull electrons towards itself within a chemical bond, increasing across and up the periodic table.

Fluorine

The most electronegative element on the periodic table, located at the top right (excluding noble gases).

Polar Bond

A covalent bond where electrons are shared unequally because one atom has higher electronegativity, creating partial positive ($\delta^+$) and partial negative ($\delta^-$) regions.

Nonpolar Bond

A covalent bond with equal electron sharing; specified for the exam as occurring strictly when an element is bonded to itself (e.g., $Cl-Cl$, $H-H$, $O-O$, $N-N$) or between carbon and hydrogen ($C-H$).

Dipole

A separation of electrical charge in an unequally shared polar bond, indicated by partial charge symbols ($\delta^+$, $\delta^-$) or an arrow pointing toward the more electronegative atom with a plus sign at the tail.

Valence Electrons

The outermost electrons counted by an element's column in the periodic table across the s and p blocks (excluding transition metals in the d and f blocks).

Core Electrons

Inner-shell electrons that require a drastically higher ionization energy to remove than outer valence electrons.

Lewis Dot Structure

A diagram representing the valence electrons of molecules and ions using lines for bonding electron pairs (each representing two electrons) and dots for lone pairs.

Central Atom (Lewis Rules)

The atom placed in the middle of a Lewis dot structure, which must be the least electronegative element; hydrogen is never placed in the center because it can only hold two electrons.

Bracket Notation (Lewis Structures)

Square brackets drawn around a complete Lewis dot structure with the overall charge indicated outside to account for extra or missing electrons in a polyatomic ion.

Resonance

The ability to draw two or more valid Lewis dot structures for the same molecule by altering electron/bond placements without changing the positions of the atoms.

Resonance Hybrid

The actual physical state of a molecule displaying resonance, representing a simultaneous mixture or average of all valid contributing Lewis structures rather than oscillating between them.

Formal Charge

The hypothetical charge calculated for an atom in a Lewis structure: $\text{Formal Charge} = \text{valence electrons} - \frac{1}{2}(\text{shared electrons}) - \text{unshared electrons}$, or simplified as $\text{valence electrons} - (\text{lines} + \text{dots})$.

Optimal Formal Charge Criteria

Rules used to determine the best resonance contributor: minimize overall formal charges toward zero, and place any unavoidable negative formal charge on the most electronegative element.

Incomplete Octet

An exception to the octet rule where stable molecules contain elements with fewer than eight valence electrons: hydrogen (holds 2), beryllium (holds 4), and boron/aluminum (commonly hold 6).

Radical

A molecule that contains an odd total number of valence electrons, resulting in an unpaired electron being assigned to the least electronegative atom.

Expanded Octet

An exception to the octet rule applicable to elements in the third row and lower on the periodic table, which can accommodate more than eight valence electrons (up to 12 electrons) placed on the central atom.