Periodic properties

Pauli’s Exclusion rule

no 2 electrons can have all 4 quantum numbers be the same 

Hund’s rule

fill across degenerate energy levels before completing orbital

Electron configuration exceptions

1. if highest occupied orbital is d, remove from s first

2. half filled and full filled orbitals are more stable

Effective nuclear charge

nuclear charge felt by an electron when both the actual nuclear charge and the shielding effects of other electrons are taken into account

Atomic radius

Cations - lose electrons

• Fewer electrons than protons = smaller radius

Anions - gain electrons

• More electrons than protons = larger radius

Ionization energy

minimum energy required to remove an electron from its ground state

Exceptions: half filled p subshells (p³) and fully filled s orbitals (s²) have more

Electron affinity

ability of an atom to attract electron density in an isolated atom

Exceptions: half filled p subshells (p³) and fully filled s orbitals (s²) have less

Electronegativity

ability of an atom to attract electron density in a chemical bond

polar bond

electrons are shared unequally

Significant difference in electronegativity (further away on the periodic table)

non polar

electrons are shared equally

Little difference in electronegativity (closer on the periodic table)

Ionic

metal and nonmetal

trades electrons

covalent bond

nonmetals

shares electrons

metallic bond

metals

delocalized electrons

electrically conductive

polarizability

anion

how easily an anion's electron cloud can be shifted

1. Size of electron cloud

• Larger atomic radius = more polarizability

2. Charge and electron:proton ratio

• More negative charge = higher electron:proton ratio = more polarizability

3. Number of energy shells

• More energy shells = more polarizability

4. Effective nuclear charge

• More electrons that are further from the nucleus

• Nucleus has less pull on far off electrons

polarizing power

cation

the ability of a cation to pull back electron density in an ionic bond

1. Size of electron cloud

• Smaller atomic radius = more polarizing power

2. Electron:proton ratio

• More positive charge = higher electron:proton ratio = higher polarizing power

3. Number of energy shells

• Less electron shells = more polarizing power

4. Effective nuclear charge

• Less electrons closer to nucleus

• Nucleus has larger effect on electrons that are closer

covalent character

More polarizable anion

More polarizing power

bond energy

The process of forming bonds involves attraction (energy release) and repulsion (energy absorption)

• too close = Repulsion increases, making the system unstable.

• too far = Attraction weakens, and the bond won’t hold

Atoms naturally move toward the lowest energy state, where bonds form and the system becomes stable

• Low energy = more stable

bond energy factors

1. Number of electrons causing the bonding to occur

• More electrons = stronger bond

2. Size of atom

• Smaller atoms = stronger bonds

• Electron clouds overlap more effectively, allowing for a stronger attraction between the nuclei and the shared electrons