Bonding and chemical interactions

ionic bonding

between metal and nonmetal

one or more electrons are transferred from an atom with a low ionization energy to an atom with a high electron affinity

covalent bonding

electron pair is shared between two atoms; nonmetals that have relatively similar electronegativities

ionic bonds

atom that loses an electron becomes a cation

atom that gains an electron becomes a anion

ionic bond is a result of the electrostatic force of attraction between the opposite charges of these ions

Properties

• high melting and boiling points

• dissolve readily in water and other polar solvents

• good conductors of electricity in aqueous state

• form a crystalline lattice in solid state

covalent bonds

contain discrete molecular units with relatively weak intermolecular interactions

• low melting and boiling points

• poor conductors of electricity in liquid or aqueous states

bond order

number of shared electron pairs between two atoms

bond length

average distance between two nuclei atoms in a bond

inversely proportional to bond strength

• more bonds = smaller length

  • 1 > 2 > 3

bond energy

energy required to break a bond by separating its components into their isolated, gaseous atomic states

relative bond strength

• 3 > 2 > 1

polarity

when two atoms have a relative difference in electronegativity values

creates a dipole with the positive end of the dipole at the less electronegative atom and the negative end at the more electronegative atom

diatomic molecules

H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

polar covalent bond

uneven sharing of electrons creating electron densities with a positive and negative end

dipole moment

vector quantity of a polar bond or polar molecule

p = qd

q = magnitude of charge

d = displacement vector separating the two partial charges

formal charge

formal charge = V - N_nonbonding - ½ N_bonding

V = valence electrons

N_nonbonding = # of bonding electrons (double the # of bonds)

N_unbonding = # of nonbonding electron

sigma bond (σ)

allow for free rotation about their axes because the electron density of the bonding orbital is a single linear accumulation between the atomic nuclei

pi bond (π)

do not allow for free rotation because the electron densities of the orbitals are parallel and cannot be twisted in such a way that allows for continuous overlapping of the clouds of electron densities

hydrogen bonds

HF > HO > HN