1. Atomic Structure Content Shells, subshells & orbitals
What is an orbital?
Aufbau principle
Electron configurations
Spin & electrons in boxes
Ionising transition metals
Successive ionisations
Periodic trends
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2. Shells, subshells and orbitals
The subshells are s, p, d and f
The number of subshells in a shell is equal to the shell number:
First shell: 1s
Second shell: 2s, 2p
Third shell: 3s, 3p, 3d
Fourth shell: 4s, 4p, 4d, 4f
There is ONE orbital in the s subshell, THREE in the p and FIVE in the d subshell.
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3. What is an orbital? An orbital is a region of space described by the
wavefunction (found from the square modulus of ψ).
They are defined so that there is a 90% probability of finding the electron in the orbital.
ψ is a mathematical function. It is derived as a solution to the Schrödinger equation.

4. Shapes of s and p orbitals

5. Shapes of d orbitals

6. Subshell energies
In a one-electron atom, ie hydrogen, the subshells within a given shell have the same energy.
In a multi-electron atom, electrons shield each other. Since electron shielding depends on the shape of the orbitals, subshells within a shell are affected differently by shielding and so have different energies.
The ordering of subshell energies in an atom is given by the Aufbau principle.

7. Shielding The radial distribution function on the right shows how 1s
electrons shield 2s electrons.

8. Aufbau principle
The shells fill up in the order from following the diagonal
arrow in the diagram:
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This gives the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, etc

9. Electron configurations
These are written in the order of subshell energy, with
the number of electrons in each subshell written as a
superscript. eg Lithium is 1s2 2s1.
Electrons are commonly abbreviated to just the outer electrons
that surround the nearest noble gas configuration (which is
given in square brackets).
eg Potassium may be written [Ar] 4s1.

10. Electron spin
Electrons possess a mysterious property called spin. This is
intrinsic to the electron, ie regardless of which orbital it
occupies. A classical analogy would be a particle spinning on
its own axis. Electrons may take one of two spin states, known
as ‘up’ and ‘down’. These are depicted with up or down arrows.
A classical analogy might be the electron spinning either
clockwise or anti-clockwise on its own axis.
An electron orbital has a maximum occupancy of two electrons,
as long they have opposite spins. This is the Pauli principle.

11. Electrons-in-boxes notation
This notation shows how the electrons fill the individual orbitals.
The spin of the electron is shown in each orbital as an up- or
down-pointing arrow.
When the number of electrons is equal to or less than the
number of orbitals, then the electrons occupy separate orbitals.
These electrons are shown all pointing up.
Only when the number of electrons exceeds the number of
orbitals are they shown as paired.

12. Examples
Carbon 1s2 2s2 2p2
Fluorine 1s2 2s2 2p5

13. Transition metals
In transition metals the 3d and 4s subshells are very close in
energy, which leads to some anomalous behaviour...
Because of the stability of the half-full and totally-full 3d
subshell there are two exceptions to the Aufbau principle:
Cr is [Ar] 4s1 3d5 and Cu is [Ar] 4s1 3d10. These are the
ONLY exceptions in the first row d block.
2. The 4s electrons are ionised before the 3d. This is
because the ordering of 3d and 4s orbital energies is
reversed in cations.
eg Fe2+ is [Ar] 3d6 NOT [Ar] 4s2 3d4.

14. Successive ionisation energies
They are defined such that one electron is removed at a time.
These ionisations are always for the gas phase.
1st Ionisation energy: M(g) → M+(g)
2nd Ionisation energy: M+(g) → M2+(g)
Successive ionisation energies for an element always
increase. This is because the negative electron has to be
removed from an increasingly positive ion.
The jump between successive ionisation energies is particularly
large after a noble gas configuration is reached.

15. Successive ionisation energy graph
Note how jumps between subshells and in going to unpaired
electrons are insignificant compared to the noble gas jump.

16. Periodicity: 1st ionisation energy

17. Explanation of ionisation energy trend
General increase – due to increase in nuclear charge
while electrons are added to the same shell (which means
little increase in shielding across the period).
Drop between Mg and Al – as 3p subshell is higher in
energy than 3s, so less energy needed to eject the electron.
Drop between P and S – as electrons start pairing when
four electrons are in a p subshell. This pairing increases
the energy of the electrons, so they are easier to ionise.

18. Periodicity: radii
Atomic radii DECREASE across a period due to the increasing nuclear charge with little increase in shielding. Positive ions are smaller than their respective neutral atoms, due to the increasing proton:electron ratio, especially if the all the valence electrons are ionised. Negative ions are larger than their respective neutral atoms, due to the decreasing proton:electron ratio and the extra electron-electron repulsion from the additional electron.

19. Graph of ionic radii across period 3

20. Electronegativity
Generally defined as the nuclear attraction experienced by the valence (ie outer shell) electrons. The Pauling definition is: the attraction of the bonding electrons in an atom to the nucleus. It is useful for predicting bond dipoles. Electronegativity increases from left to right across a period, and up all groups. Fluorine regarded as the most electronegative element. Caesium is the least electronegative (or most electropositive) element. Increasing ionisation energy and decreasing atomic radii increase electronegativity.

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