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Trends in Atomic Properties Atomic Radius There are various definitions of the size of an atom.

van der Waals radius: half the distance of closest contact between two atoms of an element without them bonding – much longer than bonding radii

Metallic radius: half the distance between two nuclei of a metallic element in the solid metal

Covalent radius: half the distance between two nuclei in a non-metallic element when connected by a single bond

Metallic r (of Ms) and covalent r (of NMs) are generally called atomic radii General trend (best for s and p blocks): r across row (as Zeff )

r down a column (as valence n )

Ionization Energy (IE) ΔH for reaction A(g) A+

(g) + e–

always a (+) value, always unfavourable

measures energy of highest-energy occupied orbital: IE = –E(HOMO) General trend: IE across row (as Zeff ),

IE down a column (as valence n ) But, small discontinuities due to electronic configuration

Electron Affinity (EA) ΔH for reaction A(g) + e– A–

(g)

(–) value if e– is favoured, (+) value if disfavoured

measures energy of lowest-energy unoccupied orbital: EA = E(LUMO) Most favoured: Gr 16, 17 Most disfavoured: Gr 2, Gr 18 Note: EA of A = –IE of A–

F + e– F– ΔH = –3.4eV, so EA(F) = –3.4eV

F– F + e– ΔH = +3.4eV, so IE(F–) = +3.4eV

0

5

10

15

20

25

1 2 13 14 15 16 17 18

Ion

izati

on

En

erg

y /

eV

Group Number

n = 1

n = 2

n = 3

n = 4

Li

F

Na

Cl

K

Br

50

70

90

110

130

150

170

190

210

230

0 10 20 30 40

rad

ius /

pm

atomic number

Atomic Radii

ionization energy:

A A+ always (+)

electron affinity:

A A– note the sign! (–) if favoured

(+) if disfavoured

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Electronegativity (χ) a variable concept, with multiple formal definitions

conceptually: tendency of an atom to attract/hold electrons

general trend: χ across row (as Zeff ), χ down a column (as n ) Pauling Electronegativity (χP)

measures ability of an atom to attract shared electrons in a bond

atomic parameter, but only meaningful in context of molecules with covalent bonds

used to explain bond and molecular dipole moments (δ+ and δ– on atoms)

defined in terms of covalent bond strengths as a basis for valence bond theory (see later): scale from 0.8 (Cs) to 4.0 (F)

Mullikan or Absolute Electronegativity (χM)

based on IE and EA, purely atomic parameter

conceptually, average of IE and EA values: χM = ½(IE – EA)

scaled to Pauling by χM = 1.35[½(IE – EA)]½ – 1.37, with IE and EA measured in eV

Atomic Properties: Summary

Across a row: Zeff , so orbital size / energy , so: r , IE , EA more favoured, χ

Down a column, n so orbital size / energy , so: r , IE , EA less favoured, χ

Trends are rationalized in terms of orbital energies/sizes, controlled by Zeff and n Discontinuities are best explained in terms of specific electronic configurations

IE & EA measure orbital energies. IE & EA trends (and trend discontinuities) are experimental evidence that SWE-predicted orbital energies are correct.

Most important properties for an understanding of bonds and molecules: orbital energy and electronegativity

E

0

M

Electronegativity

M = ½(IE – EA)

EA

IE

small r high IE high χ

(favourable EA)

large r low IE low χ

(unfavourable EA)