CHM 401


Periodic Properties


Consider Li:

How good is such an approximation?

The reduction of nuclear charge seen by outer electrons due to the inner electrons is called screening or shielding


Z* = Z - σ

Z* = effective nuclear charge

Z = true nuclear charge

σ = shielding constant

Shielding constants are established by ad hoc rules or quantum mechanical calculations

Trends in Z*

s orbitals are good screeners, p orbitals are moderate screeners; d, f orbitals are very poor screeners.

This means changes across the Table are bigger than changes down the Table


Experimental Properties

Ionization Potential (IP) or Ionization Energy (IE)

A(g) A+(g) + e

The energy associated with this reaction is the ionization energy

IP or IE is thermodynamically positive

IP generally follows Z* across the Table


Electron Affinity : EA

A(g) + eA(g)

EA reported with the wrong thermodynamic sign

Periodic trends are reverse of IP

Atomic Size : Radius of an atom - a difficult concept to define

atomic radius depends upon the bonding situation

metallic radius : ½ the internuclear distance between atoms in the metallic state

covalent radius : ½ the internuclear distance in a homonuclear covalent bond

ionic radius : size of an ion in a solid

van der Waals radius : the distance at which an atom just starts to respond to interactions from neighboring atoms

radii all of different numerical values but similar periodic trends:

r increases as go down the Periodic Table (bigger n)

r decreases as go across the Periodic Table (bigger Z*)


Electronegativity generally used as a qualitative concept

Pauling defined electronegativity as "the power of an atom to attract electrons to itself in a bond"

Pauling Scale: where:

Δ = E(A-B) – [E(A-A) + E(B-B)]/2

E(A-A), E(B-B), E(A-B) are bond energies in kJ/mol

Allred-Rochow scale: suggested that electronegativity is an electrical force


rcov is covalent radius in Angstroms

Mulliken: made electronegativity into a true atomic property; now called absolute EN