Chapter 4: Chemical Bonding and Molecular Geometry

Ionic bonding

attraction between the opposite charges of a cation and an anion, complete transfer of electrons

Lattice energy

the energy required to completely separate a mole of a solid ionic compound into its gaseous ions

Coulomb’s Law

force between to particles based on their charge and size

F= k(Q1Q2/r²)

Covalent bond

bond in which two electrons are shared by two atoms

Bond enthalpy

change in enthalpy required to break a bond

Covalent compounds

relatively weak intermolecular forces, liquid, gases, low melting solids

Ionic compounds

stronger forces, solids with high melting points

Non-polar covalent bonds

electrons shared equally

Polar covalent bonds

electrons not shared equally, one atom pulls more

Electronegativity (EN)

the ability of an atom to attract electrons to itself (determines whether something is polar or non-polar)

Bond polarity

small EN difference = non-polar covalent

mid EN difference = polar covalent 

Naming binary molecules

2 non-metal elements in the formula, 2nd element gets “ide” ending, greek prefixes indicate number of atoms

Naming acids

1. anions ending in “ide” get hydro prefix, “ic” ending and add “acid” to name

2. anions ending in “ate” get “ic”, ending in “ite” get “ous” ending

Naming hydrates

greek prefixes to indicated number of water molecules

Octet rule

atoms, other than hydrogen, tend to form bonds until surrounded by 8 electrons (have a full octet)