Chemical Bonding Flashcards

Electrostatic attraction between oppositely charged ions, formed when electrons are transferred from a metal to a nonmetal.

A bond formed by the sharing of one or more electron pairs between two nonmetal atoms.

Bonding in metals where valence electrons are delocalized in a 'sea of electrons' shared among a lattice of metal cations.

The group number (1A-8A) equals the number of valence electrons. For example, oxygen (Group 6A) has 6 valence electrons.

Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, mimicking a noble gas configuration.

Count total valence electrons, place the least electronegative atom in the center, distribute electrons as bonding pairs then lone pairs, and satisfy octets.

Formal charge = valence electrons - lone pair electrons - (1/2 bonding electrons). The best Lewis structure minimizes formal charges.

Two or more valid Lewis structures for the same molecule that differ in electron placement. The true structure is a hybrid of all resonance forms.

Predicting molecular geometry by minimizing repulsion between electron domains (bonding and lone pairs) around a central atom.

Tetrahedral with bond angles of approximately 109.5 degrees. Four bonding pairs, zero lone pairs.

Trigonal pyramidal (e.g., NH3). The electron-domain geometry is tetrahedral, but the lone pair compresses bond angles to about 107 degrees.

Bent (e.g., H2O). Bond angle is approximately 104.5 degrees due to lone-pair repulsion.

Trigonal planar (3 bonding domains, 0 lone pairs), such as BF3.

A molecule is polar if it has polar bonds and an asymmetric geometry so that dipole moments do not cancel.

Its linear geometry causes the two equal bond dipoles to point in opposite directions and cancel exactly.

Weak intermolecular forces caused by temporary dipoles from electron fluctuations. Present in all molecules; strength increases with molar mass and surface area.

Attractive forces between the positive end of one polar molecule and the negative end of another. Stronger than London forces alone for molecules of similar size.

A strong dipole-dipole interaction where H bonded to F, O, or N is attracted to a lone pair on F, O, or N of another molecule.

Higher bond order means stronger and shorter bonds. A triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond.

A sigma bond is head-on overlap along the internuclear axis (first bond). A pi bond is side-by-side p-orbital overlap (second and third bonds in double/triple bonds).

The energy required to break one mole of a bond in the gas phase. It can estimate enthalpy of reaction: delta H = sum(bonds broken) - sum(bonds formed).

Stronger intermolecular forces require more energy to overcome, resulting in higher boiling points. Hydrogen bonding > dipole-dipole > London dispersion (for similar-sized molecules).