1. Electronic structure in atoms
AH Chemistry, Unit 1(a)

3. Bohr’s theory cannot explain spectra for atoms more complex than hydrogen, which show sub-lines.
In 1926 Erwin Schrodinger devised a theory to describe more complicated atoms: quantum mechanics.
This classifies regions of atoms into SHELLS, SUB-SHELLS and ORBITALS.

4. Quantum numbers
Each electron in an atom is described by 4 different quantum numbers:
n, l, ml, ms
The first three (n, l and ml) describe an atomic orbital – where an electron is found
3 numbers because three dimensions
A fourth quantum number describes the spin of an electron

5. Quantum Numbers 1 and 2 SHELLS
1. PRINCIPAL QUANTUM NUMBER (n): one on which energy of an electron principally depends; orbitals with same n are in same shell; also determines size of an orbital; values = 1, 2, 3,…

6. Quantum Numbers 1 and 2 SUB-SHELLS
2. ANGULAR MOMENTUM QUANTUM NUMBER (l): energy of an atom also depends to a small extent on l; distinguishes orbitals of same n by giving them different shapes; any integer value from 0 to n-1; orbitals of same n but different l are in different sub-shells:
s p d f g
Example: 2p indicates shell 2 (energy), sub-level p (shape)

9. Werner Heisenberg stated in 1927 that it is impossible to know with precision both the position and the momentum of an electron.
The observer affects the observed.
Only noticeable on the sub-atomic scale (e.g. an electron, not a baseball, nor dust).
It is not possible to define a point in space where an electron will be found.
However, we can obtain the probability of finding an electron at a certain point.

10. An area where there is a greater than 90% chance of finding an electron = atomic orbital

11. Quantum Number 3 ORBITALS
3. MAGNETIC QUANTUM NUMBER (ml): distinguishes orbitals in same shell and subshell (i.e. energy and shape) by giving them a different orientation in space; values = -l to +l

12. s orbitals

13. p orbitals

14. d orbitals

16. n = 3 l = 0 l = 1 l = 2 -1 -2 +1 +2 3s 3p 3d n = 2 l = 0 l = 1 -1 +1-1 -2 +1 +2 3s 3p 3d
n = 2
l = 0
l = 1 -1 +1 2s 2p
SHELL
n = 1
l = 0 1s
SUB-SHELL
ORBITAL

17. Pauli Exclusion Principle
“No two electrons in any one atom can have the same four quantum numbers”
An orbital can hold at most two electrons, and then only if the two electrons in that orbital have opposite spins

18. Quantum number 4
4. SPIN QUANTUM NUMBER (ms): each orbital can hold 2 electrons, and each has a different direction of spin, -½ and +½

20. Hund’s Rule
“Electrons fill orbitals singly and with parallel spins before pairing occurs”.

22. Electronic configurations
Orbital notation:
Spectroscopic notation:
1s2
2s2
2p1
Outer electrons, control chemical properties
[He]
2s2
2p1

23. Practise
Write electronic configurations (obital-box and spectroscopic) for:
Helium
Beryllium
Oxygen
Aluminium
Calcium
Bromine

24. Aufbau principle
The Aufbau principle is a building up principle, which helps you to write electronic configurations
“When electrons are placed in orbitals, the energy levels are filled up in order of increasing energy”

25. Filling orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

26. Energy depends on n and l
Different orbitals in the same sub-shell have the same energy
Orbitals with the same energy are called degenerate,
There is interaction among different sub-shells in higher energy levels, which lowers their energy and explains the order of filling.