CHM 501 Lecture

Ionization Potentials (IP)

A(g A+(g) + e

IP is thermodynamically positive (not favored) generally follows Z*

Electron configurations can account for some of the fine structure, but not all

Electron Affinity (EA)

A(g) + e A(g)

Usually thermodynamically negative but reported with the wrong sign (i.e., reported as a positive number for most elements)

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 have 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*)


Related Atomic Properties:

Absolute Electronegativity χ = (IP + EA)/2

Absolute Hardness η = (IP – EA)/2

(Hardness is a measure of how easily an electron moves around in space about an atom.)

Electronegativity

Pauling: "the tendency of an atom in a bond to attract electrons to itself"

Pauling's thermodynamic quantitation:

Δ = DAB -(DAA + DBB)/2

D are bond energies between A and A, B and B, or A and B

This requires that all bonds be of the same type (single, double, triple, etc); these can be hard to find for some atoms; requires an arbitrary standard (χH = 2.2) - this gives F as 4.0. Pauling was able to assign electronegativities to most of the Periodic Table: generally increases to the right and decreases down the Table

Allred and Rochow suggested that electronegativities are due to an electrostatic force:

Z* is the effective nuclear charge

rcov is the covalent radius (1/2 the homonuclear bond distance)

This better accounts for some of the subtleties of the Periodic Table but generally is pretty close to Pauling's values.