1. Chem 1001 Lecture 11 Four quantum numbers.
Recap – Last Lecture
Four quantum numbers.

2. Chem 1001 Lecture 4
Recap – Last Lecture
Each shell is divided into subshells called s, p, d, f….
There is one extra subshell for each new shell
First shell: 1s
Second shell: 2s and 2p
Third shell: 3s, 3p and 3d
Fourth shell: 4s, 4p, 4d and 4f
1st shell 1s
2nd shell 2s 2p
3rd shell 3s 3p 3d
4th shell 4s 4p 4d 4f
5th shell 5s 5p 5d 5f
etc

3. Recap – Last Lecture Electrons as waves.
Chem 1001 Lecture 11
Recap – Last Lecture
Electrons as waves.
Consequently electrons occupy a 3-D area.
Uncertainty means we have electron ‘clouds’ though often represented by a surface incorporating 90% of the electron density.

4. Chem 1001 Lecture 14
Sub-shell energy
In hydrogen the energy of the sub-shells of a given n are degenerate (of the same energy).
In all other atoms the sub-shells are of different energies. 
Dr Toby Hudson

5. Chem 1001 Lecture 14
Sub-shell energy
The order in which the sub-shells are filled becomes important with the orbital energy increasing in the order:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p …
Follow the arrows, starting at the top, to get the order of energies of the sub-shells.
Dr Toby Hudson

6. Electron configuration
Chem 1001 Lecture 14
Electron configuration
There are three rules in determining electron configuration:
Pauli exclusion principle - no two electrons can have the same four quantum numbers i.e. maximum of 2 electrons in any one orbital.
Aufbau principle - fill up low energy orbitals first before high energy ones.
Hund’s rule - orbitals with the same energy (i.e. the same sub-shell) have the maximum number of unpaired electrons.
Dr Toby Hudson

7. Electron Configuration
Chem 1001 Lecture 4
Electron Configuration
2 electrons can fit into a s subshell
6 electrons can fit into a p subshell
10 electrons can fit into a d subshell
14 electrons can fit into a f subshell
energy increases 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f
etc
energy increases

8. Electron Configurations
Chem 1001 Lecture 4
s: 2 electrons
p: 6 electrons
d: 10 electrons
Electron Configurations
Ground state (no ‘excited’ electrons)
H has 1 electron: 1s1
He has 2 electrons: 1s2
Li has 3 electrons: 1s2 2s1
Be has 4 electrons: 1s2 2s2
B has 5 electrons: 1s2 2s2 2p1
(1s now full)
(2s now full) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f
etc

9. Electron Configurations
Chem 1001 Lecture 4
s: 2 electrons
p: 6 electrons
d: 10 electrons
Electron Configurations
Period 2, 3 & 4
Ne has 10 electrons: 1s2 2s2 2p6
Na has 11 electrons: 1s2 2s2 2p6 3s or: [Ne] 3s1
Mg has 12 electrons: 1s2 2s2 2p6 3s2 or: [Ne] 3s2
Ar has 18 electrons: 1s2 2s2 2p6 3s2 3p6
K has 19 electrons: 1s2 2s2 2p6 3s2 3p6 4s1 or: [Ar] 4s1
Ca has 20 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 or: [Ar] 4s2
Notice! We have filled 4s before 3d

10. Electron Configurations
Chem 1001 Lecture 4
s: 2 electrons
p: 6 electrons
d: 10 electrons
Electron Configurations
Period 4
Ca has 20 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 or: [Ar] 4s2
Fe has 26 electrons: [Ar] 4s2 3d6
Zn has 30 electrons: [Ar] 4s2 3d10
Se has 34 electrons: [Ar] 4s2 3d10 4p4

11. Periodic Table Group 1: Group 2:
Chem 1001 Lecture 4
Periodic Table
Group 1:
Li: [He] 2s1
Na: [Ne] 3s1
K: [Ar] 4s1
Group 2:
Be: [He] 2s2
Mg: [Ne] 3s2
Ca: [Ar] 4s2
core valence
electrons electrons
Elements in the same group have the same valence electron configurations and, as a result, very similar chemical properties

12. Box representation of orbitals
Chem 1001 Lecture 14
Box representation of orbitals
A box represents an orbital and an arrow represents an electron.
Indicates occupancy of orbitals.
C: 1s2 2s2 2p2 Na: 1s2 2s2 2p6 3s1
Dr Toby Hudson

13. Ions Add or subtract electrons as appropriate.
Chem 1001 Lecture 14
Ions
Add or subtract electrons as appropriate.
Eg Na: 1s2 2s2 2p6 3s1 or [Ne] 3s1
Na+:1s2 2s2 2p6 or [Ne]
Eg S: 1s2 2s2 2p6 3s2 3p4 or [Ne] 3s2 3p4
S2-: 1s2 2s2 2p6 3s2 3p6 or [Ne] 3s2 3p6
Dr Toby Hudson

14. Chem 1001 Lecture 14
Ions
In forming a cation, electrons are always removed from the sub-shell with the highest principle quantum number first.
So for period 4 elements this means the 4s electrons are removed first.
Eg K: 1s2 2s2 2p6 3s2 3p6 4s1 or [Ar] 4s1
K+:1s2 2s2 2p6 3s2 3p or [Ar]
Eg Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6 or [Ar] 4s2 3d6
Fe2+: 1s2 2s2 2p6 3s2 3p6 4s0 3d6 or [Ar] 3d6
Fe3+: 1s2 2s2 2p6 3s2 3p6 4s0 3d5 or [Ar] 3d5
Dr Toby Hudson

15. Excited Configurations
Chem 1001 Lecture 4
Excited Configurations
Electrons can be excited to higher energy levels
Ground state Na has 1s2 2s2 2p6 3s1
One excited state of Na has 1s2 2s2 2p6 3s0 3p1
Excited state relaxes to ground state emitting yellow light
sodium line
1s2 2s2 2p6 3s0 3p1
The Sun
1s2 2s2 2p6 3s1 Na
emission spectra
sodium street light

16. Chem 1001 Lecture 4
Applications
Mendeleev constructed his periodic table guided by the chemical properties of the elements. The reason this works is that elements in the same Group have similar valence shell electron configurations. p s d
Periodic Table has blocks of elements:
“s block” (two elements wide)
“p block” (6 elements wide)
“d block” (10 elements wide)
“f block” (14 elements wide) f

17. Learning Outcomes: By the end of this lecture, you should:
be able to write out the electron configuration for atoms and ions.
recognise ground state and excited state electron configurations.
Be able to use the box notation to represent orbital occupancy.
Understand the three rules associated with determining electron configurations.
be able to complete the worksheet (if you haven’t already done so…)

18. Questions to complete for next lecture:
Chem 1001 Lecture 4
Questions to complete for next lecture:
List the sub-shells present in the third shell and indicate the maximum number of electrons each may contain.
Write the electron configuration of the following atoms or ions: N
P3- Ca Mn
What feature do all elements in the ‘p’ – block of the Periodic Table share?

19. Questions to complete for next lecture:
Chem 1001 Lecture 4
Questions to complete for next lecture:
The electron configuration of He is 1s2. Why do you think it is included in Group 18 in the Periodic Table rather than Group 2?
Which of the following are impossible electron configurations? Give your reason(s).
1s2 2p1
1s2 2s3
1s2 2s2 2p6 3d1
1s2 2s2 2p7
1s2 2s0.5