Chemistry flashcards, generated for you.

An attractive force between atoms or ions that holds them together in a compound, formed by the transfer or sharing of electrons.

The ability of an atom to attract electrons toward itself in a covalent bond, measured on the Pauling scale (0–4.0).

Calculate the electronegativity difference (ΔEN). If ΔEN > 1.7, typically ionic; if ΔEN < 0.4, nonpolar covalent; if 0.4 < ΔEN < 1.7, polar covalent.

A 2D diagram showing valence electrons (dots) and bonds (lines) around atoms, used to visualize electron distribution and bonding in molecules.

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding. Example: Carbon in CH₄ uses sp³ hybridization to form four equivalent C–H bonds.

Valence Shell Electron Pair Repulsion theory predicts molecular geometry by assuming electron pairs (bonding and lone) repel each other and position themselves to minimize repulsion.

Electron geometry accounts for all electron pairs (bonding + lone pairs); molecular geometry considers only bonding pairs. Example: H₂O has tetrahedral electron geometry but bent molecular geometry.

Resonance structures are multiple valid Lewis structures for one molecule where atoms and bonds don't actually shift. Molecules exhibit resonance when a single Lewis structure cannot accurately represent bonding; the true structure is a hybrid of all contributors.

Bond dissociation energy (BDE) is the energy required to break a bond homolytically. Higher BDE indicates stronger, less reactive bonds; lower BDE indicates weaker, more reactive bonds. ΔH_rxn = Σ(BDE reactants) − Σ(BDE products).

Oxygen has fewer electron shells and more protons pulling on valence electrons than nitrogen. This makes O–H bonds more polar (ΔEN = 1.24) than N–H bonds (ΔEN = 0.84), resulting in stronger hydrogen bonding and higher boiling points in O–H compounds.