MCAT Kaplan Chemistry ch.3 Bonding and Chemical Interactions

Octet rule

Elements will be most stable with eight valence electrons however there are many exceptions

Ionic bond

Formed via the transfer of one or more electrons from an element with a relatively low ionization energy to an element with a relatively high electron affinity

Cation

Positively charged ion

Anion

Negatively charged ion

Crystalline lattice

Large, organized array of ions in ionic compounds

Covalent bond

Formed via the sharing of electrons between two elements of similar electronegativities

Bond order

Whether a covalent bond is a single, double, or triple bond; as bond order increases, bond strength increases, bond energy increases, and bond length decreases

Nonpolar bonds

Result in molecules in which both atoms have exactly the same electronegativity

Polar bonds

Form when there is a significant difference in electronegativities, but not enough to transfer electrons and form an ionic bond; more electronegative atom takes partial negative charge

Coordinate covalent

When a single atom provides both bonding electrons while the other atom doesn't contribute any; found in lewis acid/base rxns

Lewis dot symbols

Chemical representation of an atom's valence electrons

Formal charges

Exist when an atom is surrounded by more or fewer valence electrons than it has in its normal state

VSEPR theory

Predicts the three dimensional molecular geometry of covalently bonded molecules; electrons (bonding or nonbonding) arrange themselves to be as far apart as possible from each other in 3D space, leading to characteristic geometries

Electronic geometry

The position of all electrons in a molecule whether bonding or nonbonding

Molecular geometry

The position of only the bonding pairs of electrons in a molecule

Polarity of molecules

Depends on the dipole moment of each bond and the sum of the dipole moments in a molecular structure

Sigma bonds

Result of head to head overlap

Pi bonds

Result of the overlap of two parallel electron cloud densities

Intermolecular forces

Electrostatic attractions between molecules; significantly weaker than covalent bonds which are weaker than ionic bonds

London dispersion forces

The weakest interactions; present in all atoms and molecules; as size of the atom or structure increases, so does the corresponding force

Dipole-dipole interactions

Occur between the oppositely charged ends of polar molecules; stronger than London forces; seen in solid and liquid phases but negligible in gas phase due to distance between particles

Hydrogen bonds

Specialized subset of dipole-dipole interactions involved in intra and intermolecular attraction; occurs when hydrogen is bonded to one of three very electronegative atoms (NOF)