Chemistry flashcards, generated for you.

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

The tendency of an atom to attract electrons toward itself in a chemical bond, measured on a scale (Pauling: 0.7–4.0).

An electron transfers from a metal atom (low electronegativity) to a nonmetal atom (high electronegativity), creating oppositely charged ions that attract electrostatically.

Polar covalent: unequal electron sharing due to different electronegativities (ΔEN 0.4–1.7), creating partial charges (δ+ and δ−). Nonpolar covalent: equal or nearly equal sharing (ΔEN <0.4), no significant partial charges.

A diagram showing valence electrons of atoms as dots and bonds as lines/pairs; represents electron distribution and bonding in a molecule.

Mixing of atomic orbitals to form new hybrid orbitals with different shapes/energies for bonding. Example: carbon in CH₄ undergoes sp³ hybridization, forming 4 equivalent sigma bonds.

Sigma bond: head-on overlap along internuclear axis; allows rotation; found in single and multiple bonds. Pi bond: sideways overlap of p-orbitals perpendicular to internuclear axis; prevents rotation; found only in double/triple bonds.

Electron groups (bonding + lone pairs) around a central atom repel each other, maximizing distance apart to minimize repulsion, determining the 3D shape of the molecule.

The individual bond dipoles cancel due to molecular symmetry (geometry). Example: CO₂ is linear with two equal, opposite C=O dipoles that vectorially sum to zero.

Water has two lone pairs on oxygen (high electronegativity, EN=3.44) and two polar O-H bonds, enabling two H-bonds as acceptor and two as donor, creating a more extensive hydrogen-bonding network than NH₃ (one lone pair, three N-H bonds).