2c - Periodic Properties

Effective Nuclear Charge

Z_eff is the net positive charge experienced by a valence e^-, accounting for both nuclear attraction and shielding.

⬆Z_eff= less shielding = more nuclear attraction

⬇Z_eff= more shielding = less nuclear attraction

Periodic trend of Z_eff

✦ Increases across a period because nuclear charge Z increases while same-shell e^- only partially shield

✧ Slightly increases down a group, but additional core e^- shells significantly increase shielding (less nuclear attraction)

Periodic trend of Ionization Energy

Ionization energy (IE) = the energy required to remove an e^- from an atom

✧ Decreases down a group because the e^- being removed successively larger shells.

✦ Increases across a period because of incomplete screening from other e^-

Ionization energy trend exceptions

✦ discontinuity between ns²np⁰ & ns²np¹: 2s penetrates closer than 2p = 2s in lower in E = need more E to eject e^- form 2s

✦ discontinuity between np³ & np⁴ (adding a 4^th e^- to a ½-filled shell): maximizing exchange E 𝚷_e while relieving pairing E 𝚷_c

∆E

The stabilizing energy gained when electrons with parallel spins occupy degenerate orbitals. The lower the value, the more stable

𝚷_c

Energy required to pair two electrons in the same orbital, overcoming their repulsion. Less pairing E increases stability

𝚷_e

Exchange energy. Stabilizing energy gained when electrons with parallel spins occupy degenerate orbitals. This lowers the total energy ∆E and increases stability.

Periodic trend of Electron Affinity

EA = the energy of adding of an electron to a neg. ion

✧ Increases (more -) across a period due to increasing Z_eff and stronger nuclear attraction

✧ Decreases (less -) down a group due to increased atomic size and shielding = e^- addition less favoured

**less E is needed to remove e^- from an anion because of higher shielding

Periodic trend of Electronegativity

EN - the tendency of an atom to attract bonding e^-

✦ Increases across a period due to higher Z_^eff (e^- pulled together = smaller atomic radius)

✧ Decreases down a group due to increased atomic size (e^- are farther from nucleus) and shielding. This decrease is more marked in groups on the right side

**𝛘 of noble gases is “irrelevant” as they aren’t very reactive

Electronegativity scales

✧ Mulliken Scale - atom has high 𝛘 if it has high IE and/or high +EA

✧ Allen Scale - designed for the main group elements, uses ionization enthalpy data (∆H_IE)

✧ Allred Rochow Scale - considers the Z_eff experienced by a bonding e^-.

✧ Pauling Scale - considers the difference in bond strength between a bond’s real strength and its calculated “non-polar” strength

Periodic trend of Atomic Radii

✦ Atomic radius increases down a group as new e^- shells are added

✧ Decreases across a period due to increasing Z_eff (⬆Z and only partial shielding by same ns,np e^-)

Coordination number

# of attachments between a metal and its ligands

Ionic Radii

Ions get larger going down a group, with increasing coordination number

✦ cations are smaller than parent ions: fewer e^-, higher Z_eff, e^- are held tighter

✦ anions are larger than parent atoms: more e^-, lower Z_eff, e^- held looser

✧ isoelectronic species decrease in size as Z increases: same shielding but higher Z = higher Z_eff = e^- are pulled closer (O^2->F^->Na⁺>Mg²⁺)

Metallic radius v.s. Covalent radius

Metallic radius: ½ the distance between the nuclei of metal atoms in the solid state

Covalent radius: ½ the distance between identical atoms in a covalent compound

These 2 radii = atomic radii

Lanthanide Contraction

The gradual decrease in atomic radii across the lanth. series due to poor shielding by 4f e^-

✦ similar sizes between 6th and 5th row transition metals

✦ 6th row d-block transition metals are smaller than “expected”