Periodic Properties of the Elements

Chapter 7

Effective Nuclear Charge

Orbitals of the same energy are said to be degenerate.

Effective nuclear charge is the charge experienced by an electron on a many-electron atom.

The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons.

Effective Nuclear Charge

Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nucleus.

The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons.

As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.

As the distance from the nucleus increases, S increases and Zeff decreases.

Energies of Orbitals

The result of the Effective nuclear charge on the electronic configuration is a shift of the orbital ordering for large (n = 4 or more) electronic systems.

Energies of Orbitals

Sizes of Atoms

Electron Shells in AtomsConsider a simple diatomic molecule.The distance between the two nuclei is called the bond distance.If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.

Atomic Radii•As a consequence of the ordering in the periodic table, properties of elements vary periodically.•Atomic size varies consistently through the periodic table.•As we move down a group, the atoms become larger.•As we move across a period, atoms become smaller.

There are two factors at work:•principal quantum number, n, and•the effective nuclear charge, Zeff.

Atomic Radii decreasing across a period is due to:

shielding or screening effect inner electrons [He] or [Ne], etc. block the nuclear charge for 2 or 10 or

__ electrons consequently the outer electrons feel

a stronger effective nuclear charge Li [He] shields effective charge is +1 Be [He] shields effective charge is +2

Atomic Radii

Atomic Radii

All radii are in angstroms.

Ionic Radii

Cations (+ ions) are always smaller than their neutral atoms

Li Be1.52 1.11

Li+ Be2+0.6 0.31

Na Mg Al1.86 1.6 1.43

Na+ Mg2+ Al3+0.95 0.65 0.5

Ionic Radii Anions (- ions) are always larger

than their neutral atomsN O F

0.7 0.66 0.64

N3-

O2-

F-

1.71 1.4 1.36

Ionization Energy minimum amount of energy required to

remove the most loosely held electron from an isolated gaseous atom measure of an element’s ability to form

positive ions

Mg(g) + 738kJ/mol Mg+ + e-

Atom(g) + energy ion+(g)

+ e-

first ionization energyfirst ionization energy

Ionization Energy second ionization energysecond ionization energy

energy required to remove a 2nd electron

ion+ + energy ion2+ + e-

Mg+ + 1451 kJ/mol Mg2+ + e-

• can have 3rd, 4th, etc. ionization energies

Ionization Energy generally increases as you go

across a period important exceptions at Be & Mg, N & P

generally decreases as you go down a group

Ionization Energy (kJ/mol) vs Atomic Number

H1

He2

B5

C6

N7

O8

F9

Ne10

Mg12 Al

13

Si14

P15

S16

Cl17

Ar18

Ca20

Be4

Li3

Na11 K

19

0

500

1000

1500

2000

2500

0 5 10 15 20

Ionization Energy First 4 ionization Energies (kJ/mol) -

Period 3

IA IIA IIIA IVANa Mg Al Si

1stIE 496 738 578 7862ndIE 4562 1451 1817 15773rdIE 6912 7733 2745 32324thIE 9540 10,550 11,580 4356

Ionization Energy

these energies are exactly why these ions form

Na becomes Na+

Mg becomes Mg2+

Al becomes Al3+

Si does not form simple ions

Ionization Energy

Electronegativity

The attraction of an atom to electrons

This is not a measurable property BUT it is very useful in helping to predict bonding (attraction of electrons)

Electronegativity