Topic 12

Atomic Structure

HL

12.1 Electron Configuration

• Studies of the Ionisations Energies going across a period and down a group– IE- The energy required to remove an elektron

from an atom or ion

• Orbitals

+ +Nucleus

K L M N

-

Electron

Energy

1st IE: A(g) + energy A(g)+ + e-

(g) 2nd IE: A(g)

+ + energy A(g)2+ + e-

(g)

Within an energy level (shell or orbital): increased nuclear attraction => Increase in IE

If removed from inner filled levels: electrons experience a much higher effective nuclear charge => large raise in ionisation energy.

Successive ionisation energies

Variation of I.E. across a period:• Going across a period the I.E. increase. Increased charge of the

nucleus increases the effective nuclear charge and the removed electrons are on the same energy level.

Orbitals

• The orbitals are solutions to the wave-equation. They depict the probability of where the electron can be within 90 % probability

• Each orbital contains maximum 2 electrons

• The different orbital shapes are called s, p, d, f, g.

S-orbital: sphere formelectrons fill up S-orbitals in group 1 and 2

K-shell: 1s , L-shell: 2s , M-shell: 3s

p-orbitals dumbbell formThree at each energy level except level 1. They are called px, py, pz and are orientated in three different directions in space

electrons fill up p-orbitals in groups 13-18L-shell: 2p , L-shell: 3p

d-orbitalsFive at each energy level except level 1 and 2. They are orientated in three different directions in space.

Electrons fill up d-orbitals in group 4-12 (transition elements)M-shell: 3d , N-shell: 4d , O-shell: 5d

f-orbitalsSeven at each energy level except level 1, 2 and 3. They are orientated in three different directions in space.

Electrons fill up f-orbitals for the lanthanoids and actiniodsN-shell: 4f , O-shell: 5f

• Each orbital can just be occupied by two electrons. And the electrons must have different spin (Pauli principle)

• The relative energies: s<p<d<f

Shells Energy level Orbital’s Number of e-

K n=1 1s 2

L n=2 2s2p 8

M n=3 3s3p3d 18

N n= 4 4s4p4d4f 32

The Aufbau principle

• One s-orbital on an energy level• Three p-orbital’s on an energy level• Five d-orbital’s on an energy level• Seven f-orbital’s on an energy level

• Electrons always try to adopt the lowest energy configuration as possible. This is known as the “Aufbau principle”

• The orbital in the different energy levels overlaps each other- “read” the periodic table

• You should be able to apply the Aufbau principle for an atom up to Z = 54.

Hund’s rule

• When orbitals of identical energy are available, electrons occupy these singly rather than in pairs.

Electronic configuration of ;

• 6C: 1s22s22p2

• 16S: 1s22s22p63s23p4

• The electronic configuration of atoms and their ions can be explained in a good way with the Aufbau principle.

Formation of ions

• Atoms like to have filled orbitals, almost as good as that is to have half filled orbitals.

• Sometimes, i.e. Zn Zn2+, the atom throw away electrons on a lower level, 4s2, instead of breaking the full 3d10 on the higher energy level => all orbitals are filled.

• Also for Cr and Cu => 4s2 => 4s1 and e- to d-orbitals to make half-filled and filled orbitals

Back to variation of I.E. across a period:• Going across a period the I.E. increase. Increased charge of the nucleus

increases the effective nuclear charge and the removed electrons are on the same energy level.

But:• From second to third element decrease: Electron removed from p-

orbital instead of s-orbital.• From fifth to sixth element decrease: In fifth orbital's are singly

filled, in sixth one is doubly filled => higher repulsion.

If removed from inner filled level: electrons experience a much higher effective nuclear charge => large raise in ionisation energy.