1. Electrons!!
You’ll be happy to know that we are leaving calculations for a little while…
In small groups (max 5) brainstorm what you know about the arrangement of electrons in an atom/ion

2. Electron configurations…
A few useful terms i.e. LEARN THEM
Quantum Number – refers to the electron shell number (or Period if you want to be lazy)
n – is the symbol for the Quantum Number (don’t confuse it with the n we use for moles)

3. Electron configurations…
The quantum number can be used to work out the maximum number of electrons in that energy level.
We use:
Number electrons = 2n2
Predict the maximum number of electrons in shells 1 and 3

4. Electron configurations…
Electrons fill the orbitals in pairs. Why don’t they repel each other as the a both NEGATIVELY charged?
Spin

5. Electron configurations…
Electrons have a property called spin.
They are either
Spin Up or
Spin Down

6. Electron configurations…
In order for electrons to pair they must have opposite spins. Either Up/Down or Down/Up (yes there is a difference but we can ignore it for now)
Electron spin is the basis for MRI techniques and NMR – which we will have a play with in the summer term

7. Electron configurations…
Always fill the up arrow first
Always entirely fill a section before moving on to the next one
In a section with more than one box fill ALL up arrows before starting on the down ones

8. Electron configurations…
Now to explain what the boxes are.
The planetary model of the atom was disproved and replaced with orbital theory.
Orbital theory states that there are different energy levels in an atom and electrons are restricted to one of these “quanta”

9. Electron configurations…
S-orbitals
These are the lowest energy orbital in a shell
There is ONE s-orbital in every quantum number ≥ 1
Each orbital contains a max of 2 electrons
They are spherical

10. Electron configurations…
P-orbitals
These are higher in energy than s-orbitals
There are THREE p-orbitals in every quantum number ≥ 2
Each orbital contains a max of 2 electrons
They are dumbbell shaped

11. Electron configurations…
D-orbitals
These are the second highest energy orbitals in a shell
There are FIVE d-orbitals in every quantum number ≥ 3
Each orbital contains a max of 2 electrons
They are all manner of odd shapes

12. Electron configurations…
D-orbitals

13. Electron configurations…
F-orbitals
These are the highest energy orbitals in a shell
There are SEVEN f-orbitals in every quantum number ≥ 4
Each orbital contains a max of 2 electrons
They are weirder than d-orbitals

14. Electron configurations…
F-orbitals

15. Electron configurations…
Luckily we don’t have to use box notation all the time. We more commonly give the Electronic Configuration of an atom or ion
e.g. Nitrogen – 7 electrons
1s2 2s p3

16. Electron configurations…
Rules for electron configurations:
The letters are always LOWER CASE
The number of electrons in an orbital is given by a SUPERSCRIPT number
1s2 2s p5

17. Electron configurations…
For ions it’s not much different. You use the following steps…
Work out the electronic configuration of the ATOM
Work out the charge on the ion
Add or subtract electrons as necessary using the “Last in, First out” rule

18. Electron configurations… e.g. Sodium ion
Work out the electronic configuration of the ATOM
Work out the charge on the ion
Subtract electrons
Na has 11 electrons 1s2 2s2 2p6 3s1
Na has +1 charge so loses 1 electron
Na+ has 10 electrons 1s2 2s2 2p6

19. Electron configurations…
Oh, if life was always this simple. Unfortunately once the quantum number of an element/ion is ≥ 3 the rule for filling and losing electrons gets more complicated.
e.g. Vanadium
1s2 2s2 2p6 3s2 3p6 4s2 3d3

20. Electron configurations…
The 4s orbital is lower in energy than the 3d orbital so it fills up first.
It also loses its electrons before the 3d orbital because it is further from the nucleus.
e.g. Vanadium (II)
1s2 2s2 2p6 3s2 3p6 3d3

21. Electron configurations…
This luckily only affects transition metals.
Unfortunately there are exceptions. This is due to the electronic stability of the atom/ion.

22. Electron configurations…
What you’d expect:
Chromium
Copper

23. Electron configurations…
What you actually get:
Chromium
Copper
Why are these more stable?

24. Electron configurations…
Chromium
A half filled s-orbital with 5 half filled d-orbitals is more stable than a full s-orbital with 4 half filled d-orbitals and 1 empty one.

25. Electron configurations…
Copper
A half filled s-orbital with 5 fully filled d-orbitals is MUCH more stable than a full s-orbital with 4 full and 1 half filled d-orbitals.

26. Periodicity…
You should have noticed over the past 6 years of science lessons that there are patterns to behaviour as you go down a group and across a period (if not, shame on you!)
We are now going to concern ourselves with the patterns in IONISATION ENERGIES

27. Periodicity…
First some definitions:
First Ionisation Energy
The energy required to remove ONE electron from ONE MOLE of GASEOUS atoms to form ONE MOLE of 1+ ions
e.g. Mg(g)  Mg+(g) + e-

28. Periodicity…
Successive Ionisation Energy
The energy required to remove ONE electron from ONE MOLE of GASEOUS ions (Mn+) to form ONE MOLE of ions with n+1 charge
e.g. Mg+(g)  Mg2+(g) + e-

29. Periodicity…
Atomic Radii
The distance between the centre of the nucleus and the outermost electron shell
This is using given in the unit Angstroms (Ȧ) which is equivalent to m

30. Periodicity…
(Effective) Nuclear Charge
The attractive force of the nucleus to the outermost electrons in an atom/ion
Electron Shielding
The repulsive force between inner electrons and the outermost electron shell that negates the nuclear charge

31. Home(and the rest of class)work…
Group D – Due Wed 10.45am
Group C – Due Fri 3.00pm
Work in groups (see my lists) to write a presentation explaining the following trends:
First ionisation energy in a group
First ionisation energy in a period
Successive ionisation energies in an element
Your presentation must include an element of teaching as well as a student exercise.
I will also require a breakdown of each person’s contribution to the task.